ELEMENTS   OF   INORGANIC   CHEMISTRY 


ELEMENTS 


OF 


INORGANIC  CHEMISTRY 


BY 


HARRY  C.    JONES 

ASSOCIATE   PROFESSOR   OF   PHYSICAL   CHEMISTRY   IN   THE 
JOHNS  HOPKINS  UNIVERSITY 


gorfc 
THE   MACMILLAN   COMPANY 

LONDON:  MACMILLAN  &  CO.,  LTD. 
1903 

AH  rights  reserved 


y 


COPYRIGHT,  1908, 
BY  THE  MACMILLAN  COMPANY. 


Set  up,  electrotyped,  and  published  August,  1903. 


NorfaooD 

3.  8.  Gushing  &  Co.  -  Berwick  &  Smith  Co. 
Norwood,  Mass.,  U.S.A. 


JBelucatetr  to  ».  B.  3L 

WHO   HAS   RENDERED   VALUABLE   ASSISTANCE 

IN   PREPARING   THE    ILLUSTRATIONS 

IN   THIS   WORK 


221753 


PREFACE 

THE  condition  of  chemistry  to-day  differs  greatly  from  that 
of  twenty  years  ago.  Within  this  period  the  new  physical 
chemistry  has  come  into  existence,  and  several  generalizations 
have  already  been  reached  which  affect  fundamentally  the 
whole  science  of  chemistry.  Some  of  these  can,  and,  in 
the  opinion  of  the  author,  should  be  introduced  very  early 
in  the  study  of  the  science. 

Take  the  theory  of  electrolytic  dissociation,  which  to-day  is 
as  well  established  as  many  of  our  laws  of  nature :  It  has 
shown  us  that  it  is  the  ions  and  not  the  atoms  which  are  the 
active  agents  chemically.  If  the  student  is  taught  the  con- 
trary in  the  early  stages  of  his  work,  later  this  must  all  be 
unlearned,  and  we  know  how  difficult  it  is  to  correct  first 
impressions.  If  we  ask  the  question,  why  continue  to  teach 
chemistry  from  a  purely  atomic  standpoint  after  it  has  been 
shown  not  to  be  in  accord  with  the  facts  ?  about  the  only 
answer  is  that  it  is,  perhaps,  a  little  simpler  than  the  chemistry 
of  ions.  This  reply  must,  on  reflection,  be  regarded  as  highly 
unsatisfactory.  If  a  student  can  form  a  conception  of  an 
atom  as  the  smallest  indivisible  particle  of  matter,  would  it 
be  very  difficult  for  him  to  form  a  conception  of  an  atom 
carrying  an  electrical  charge  ?  Indeed,  it  would  be  no  more 
difficult  than  to  add  to  the  conception  of  a  piece  of  metal  the 
charge  which  it  carries  when  electrical  energy  has  been  sup- 
plied to  it,  and  we  do  not  hesitate  in  physics  to  ask  him  to 
take  this  step.  Further,  since  we  know  that  the  ion  and  not 
the  atom  or  the  molecule  is  the  factor  which  enters  into  most 
chemical  reactions,  we  should  insist  upon  it  because  it  is  true. 

vii 


viii  PREFACE 

The  effect  of  mass  on  chemical  activity  is  very  great,  and 
this  should  be  pointed  out  early  in  the  study  of  chemistry. 
The  wide-reaching  generalization  known  as  the  law  of  mass 
action  may  be  discussed  somewhat  later  when  the  metals  are 
studied. 

Another  highly  important  generalization  whose  significance 
has  come  to  be  recognized  through  physical-chemical  investi- 
gations, is  the  law  of  Faraday  as  the  basis  of  chemical  valence. 
This,  as  will  be  seen,  places  the  whole  subject  of  valence  upon 
an  exact  physical  basis,  and  gives  a  definite  significance  and 
meaning  to  what  has  hitherto  been  dealt  with  in  a  rather 
vague  manner. 

The  present  method  of  treating  the  subject  of  general 
chemistry  does  not  discard  in  the  least  the  earlier  discovered 
generalizations.  The  aim  has  been  to  add  to  these  older  gen- 
eralizations those  more  recently  brought  to  light,  and  to 
incorporate  them  as  an  essential  part  of  the  subject.  This 
necessitates  a  frequent  change,  which  is  often  fundamental,  in 
our  method  of  dealing  with  chemical  phenomena,  and  in  our 
conception  of  chemical  reactions ;  but  change  has  been  intro- 
duced only  where  recent  generalizations  and  discoveries  have 
made  it  necessary,  or  at  least  highly  desirable,  in  order  to 
interpret  the  facts  in  terms  of  what  to-day  we  have  every 
reason  to  believe  is  true. 

The  physical  properties  of  substances  are  treated  rather 
more  fully  than  in  the  earlier  text-books,  since  the  tendency 
to-day,  introduced  by  physical  chemistry,  is  to  bring  chemistry 
and  physics  more  closely  together,  rather  than  to  separate 
them  as  two  distinct  sciences.  The  physical  properties  of  a 
substance  are  taken  up  after  its  chemical  properties,  in  order 
that  the  teacher  may  conveniently  omit  more  or  less  of  this 
part  of  the  text  if  the  pupil  is  not  sufficiently  advanced  to 
properly  appreciate  its  meaning. 

The  experiments  are  placed  at  the  ends  of  chapters  rather 
than  in  the  body  of  the  text  or  at  the  end  of  the  book.  By 
this  means  the  text  is  made  continuous  and  more  readable, 


PREFACE  ix 

and  while  the  student  is  compelled  to  review  the  text  in 
connection  with  the  experiment,  he  is  not  so  liable  to  copy 
it  in  describing  what  he  has  seen  and  done.  If  the  experi- 
ments were  placed  at  the  end  of  the  work,  the  beginner  is 
liable  to  miss  the  connection  between  what  he  carries  out  in 
the  laboratory  and  what  he  reads  in  the  book. 

The  problems  are  introduced  after  the  experiments  at  the 
ends  of  the  chapters.  In  solving  these  the  student  is  com- 
pelled to  review  much  of  the  text  for  a  third  time,  and 
reviews  are  the  only  means  of  becoming  really  acquainted 
with  a  new  subject.  Portions  of  my  "Principles  of  Inorganic 
Chemistry,"  which  were  considered  to  be  adapted  to  the 
beginner,  have  been  freely  introduced  into  this  book.  The 
review  of  such  portions  in  the  larger  work  will  but  serve 
to  impress  these  fundamental  facts  upon  the  more  advanced 
student. 

In  matters  of  both  experiments  and  problems  the  teacher 
must  use  his  judgment  as  to  what  should  be  omitted,  if  omis- 
sions are  necessary.  This  will  depend  largely  upon  the  time 
at  the  disposal  of  the  class,  the  laboratory  facilities,  and  the 
state  of  development  of  the  pupils. 

One  injunction  must  be  insisted  upon  at  the  risk  of  re- 
peating what  is  already  hackneyed.  Never  allow  a  student  to 
perform  an  experiment  in  a  slip-shod  manner,  with  soiled 
apparatus  poorly  set  up.  Bad  habits  thus  early  inculcated, 
like  erroneous  conceptions,  are  lost  with  difficulty,  if  lost  at 
all;-  and  a  student  is  injured  rather  than  aided  by  such  work. 

The  author  wishes  to  express  his  indebtedness  to  a  large 
number  of  friends  for  valuable  suggestions  in  preparing  this 
book,  and  especially  to  Professor  Edmund  Kenouf,  Dr.  J.  E. 
Gilpin,  and  Dr.  F.  H.  Getman. 

HABKY  C.  JONES. 


CONTENTS 


CHAPTER  I 

PAGE 

THE  SCIENCE  or  CHEMISTRY     .         .         .         .         .         .         .         .1 

CHAPTER  II 

OXYGEN        7 

4MB 

CHAPTER  III 
HYDROGEN 19 

CHAPTER  IV 
WATER       .         .        ...        .        .        ....        .32 

CHAPTER   V 
CHLORINE  .•        .         .        .        .        -.  Y  '  .        .        .        .        .        .       43 

CHAPTER  VI 
LAWS  OF  CHEMICAL  ACTION      .         .        .        .         .         .        .         .50 

CHAPTER  VII 
OZONE         .  58 

HYDROGEN  DIOXIDE  ..........      60 

\ 

CHAPTER   VIII 
COMPOUNDS  OF  CHLORINE  WITH  HYDROGEN  AND  OXYGEN       .        .       62 

CHAPTER  IX 
NITROGEN  ............       76 

COMPOUND  OF  NITROGEN  WITH  HYDROGEN 77 

COMPOUNDS  OF  NITROGEN  WITH  OXYGEN  AND  HYDROGEN       .         .       80 

xi 


xii  CONTENTS 

CHAPTER   X 

PAGE 

»    NEUTRALIZATION  OF  ACIDS  AND  BASES     ......       98 

CHAPTER   XI 

*•  THE   ATMOSPHERIC   AIR    AND   CERTAIN    RARE    ELEMENTS   OCCUR- 
RING   IN    IT    .  .  .  .  .  .  .  •  •  •  •       107 

CHAPTER  XII 
DETERMINATION  OP  RELATIVE  ATOMIC  WEIGHTS     ....     115 

CHAPTER   XIII 
THE  PERIODIC  SYSTEM 126 

CHAPTER  XIV 
BROMINE,  IODINE,  FLUORINE    ........     133 

CHAPTER  XV 
SULPHUR,  SELENIUM,  TELLURIUM      .......     150 

CHAPTER   XVI 

•PHOSPHORUS,  ARSENIC,  ANTIMONY,  BISMUTH 172 

s  COMPOUNDS  OF  PHOSPHORUS  WITH  THE  HALOGENS  ....  177 

COMPOUNDS  OF  ARSENIC  WITH  OXYGEN  AND  HYDROGEN          .         .  178 

COMPOUNDS  OF  ANTIMONY  WITH  OXYGEN  AND  HYDROGEN       .         .  180 

CHAPTER  XVII 
CARBON,  SILICON,  BORON  ........     189 

THE  ROLE  OF  CARBON  IN  PRODUCING  LIGHT 199 

CHAPTER   XVIII 

THE  METALS      .         .         .         .'. 219 

THE  ALKALI  METALS  :  LITHIUM,  SODIUM,  POTASSIUM,  RUBIDIUM, 

AND  CAESIUM 220 

SODIUM "...  220 

POTASSIUM  .  227 

AMMONIUM  232 


CONTENTS  Xlll 

CHAPTER  XIX 

PAGE 

CALCIUM,  STRONTIUM,  BARIUM 239 

CHAPTER  XX 

THE    MAGNESIUM    GROUP  :    GLUCINUM,    MAGNESIUM,    ZINC,    CAD- 
MIUM, MERCURY          .........     256 

CHAPTER   XXI 
ALUMINIUM 266 

CHAPTER   XXII 
IRON,  COBALT,  NICKEL 273 

CHAPTER   XXIII 
MANGANESE,  CHROMIUM,  MOLYBDENUM,  TUNGSTEN,  URANIUM        .     285 

CHAPTER   XXIV 
COPPER,  SILVER,  GOLD     .         .        .        .        .        .        .        .         .     296 

CHAPTER   XXV 
LEAD,  TIN       "  .        .        .        ,        .        .         .        .        ...     307 

CHAPTER   XXVI 
RUTHENIUM,  RHODIUM,   PALLADIUM,   OSMIUM,   IRIDIUM,   PLATINUM    315 

CHAPTER   XXVII 

GENERAL  RELATIONS          .........  319 

RELATIONS  WITHIN  THE  GROUPS  OF  THE  PERIODIC  SYSTEM     .         .  319 

RELATIONS  BETWEEN  THE  COMPOUNDS  or  THE  METALS  .         .         .  322 

THE  NATURE  AND  ROLE  OF  IONS  IN  CHEMISTRY      ...  325 


ELEMENTS  OF  INOKGANIC  CHEMISTEY 


CHAPTER   I 

THE   SCIENCE   OF   CHEMISTRY 

Chemistry  a  Branch  of  Natural  Science. — The  study  of  nature 
is  ever  attracting  more  and  more  attention.  We  are  brought 
in  daily  contact  with  the  world  in  which  we  live,  and  the  mind 
observes  the  processes  which  are  going  on  around  it.  These 
are  very  complex. 

If  we  turn  to  inanimate  nature,  we  see  the  rocks  crumbling 
into  dust,  and  furnishing  the  material  of  which  the  soil  is  made. 
This  is  partly  dissolved  by  the  waters,  or  carried  mechanically 
from  regions  of  higher  to  those  of  lower  level. 

If  we  turn  to  living  matter,  the  transformations  that  are 
taking  place  are  extremely  complicated.  The  seed  germinates, 
takes  root,  and  grows  at  the  expense  of  the  soil,  water,  and  air. 
The  plant  is  consumed  by  the  animal,  which  likewise  increases 
in  size.  The  animal  dies,  decays,  and  the  matter  of  which  it 
was  composed  is  given  back  again  to  the  soil,  the  water,  and 
the  air. 

The  study  of  nature,  or  natural  science,  is  so  comprehensive 
and  complex  that  it  has  been  found  necessary  to  divide  it  into 
several  branches.  This  is  especially  true  since  the  mind  is  not 
content  with  simply  observing  what  is  transpiring  before  it, 
but  places  matter  under  new  conditions  which  it  can  control  — 
experiments  with  it  —  and  observes  the  results.  The  science 
of  chemistry  is  one  of  the  branches  of  natural  science. 

The  Kind  of  Phenomena  with -which  Chemistry  is  concerned. — 
While  it  is  impossible  at 'the  outset  to  give  any  comprehensive 

" 


2  ELEMENTS  OF  INORGANIC   CHEMISTRY 

conception  of  the  science  of  chemistry,  certain  characteristics 
of  chemical  action  can  be  pointed  out. 

If  we  connect  a  piece  of  copper  with  an  electric  battery,  an 
electric  current  will  flow  through  it.  The  copper  while  carry- 
ing the  current  has  properties  which  are  different  from  copper 
through  which  no  electricity  is  passing.  Disconnect  the  copper 
from  the  electric  battery,  and  it  possesses  its  original  properties. 

Heat  a  piece  of  copper  gently,  and  some  of  its  properties  are 
changed.  It  will  give  out  heat  to  surrounding  objects ;  it  will 
occupy  a  larger  volume  when  hot  than  at  ordinary  temperatures. 
Allow  the  copper  which  has  been  warmed  to  cool,  and  it  will 
possess  again  its  original  properties ;  hammer  the  copper  or 
bend  it,  and  it  still  remains  copper.  Changes  of  this  kind  are 
known  as  physical. 

If,  on  the  other  hand,  we  heat  a  piece  of  copper  to  redness  in 
the  presence  of  the  air,  a  far  more  fundamental  change  takes 
place.  The  copper  disappears,  and  a  black  substance  is  formed 
which  has  properties  quite  different  from  the  original  copper. 
The  black  substance  does  not  look  like  a  metal.  It  cannot  be 
drawn  out  into  wire.  It  weighs  considerably  more  than  the 
original  copper,  and  in  general  has  properties  sufficiently  dif- 
ferent from  the  original  copper  to  show  that  we  are  dealing 
with  an  entirely  different  substance.  If  the  black  powder  is 
now  cooled  to  the  original  temperature  of  the  copper,  it  retains 
its  own  characteristic  properties. 

The  change  effected  in  the  latter  case  is,  then,  far  more 
fundamental  than  in  the  former.  While  certain  properties  of 
the  copper  were  changed  by  passing  an  electric  current  through 
it,  or  by  gently  warming  it,  these  properties  were  restored  again 
when  the  current  was  interrupted,  or  when  the  temperature 
was  allowed  to  fall.  The  original  copper  remained  copper. 
In  the  latter  case,  however,  the  composition  of  the  substance  was 
changed,  and  this  is  characteristic  of  chemical  activity.  The 
substances  which  react  chemically  lose  many  of  their  charac- 
teristic properties,  an4  giva  risejfc}  jieV  substances  with  very 
different  properties;  V  «  '•  •  >*•*••' 


THE   SCIENCE   OF   CHEMISTRY  3 

Elements  and  Compounds.  —  If  we  look  about  us,  we  recog- 
nize that  nature  is  made  up  apparently  of  a  great  many  sub- 
stances. The  soil  and  the  rocks  differ  greatly  in  composition 
in  different  localities,  and  are  always  more  or  less  complex. 
Water  exists  everywhere,  and  the  air  is  a  mixture  of  many 
substances.  When  we  turn  to  living  matter,  we  find  the  com- 
plexity greatly  increased.  The  simplest  living  being  is  com- 
posed of  very  complex  substances,  and  the  more  highly  developed 
organisms  contain  a  countless  number  of  substances. 

This  is  the  way  the  problem  of  the  composition  of  the  exter- 
nal world,  as  recognized  by  our  senses,  presents  itself  at  first. 
It,  however,  becomes  greatly  simplified  when  we  study  the  com- 
position of  things  in  a  systematic  and  comprehensive  manner. 
All  known  substances  fall  into  two  great  classes,  —  those  that 
cannot  be  decomposed  into  simpler  substances,  and  those  that 
can.  Take  the  piece  of  copper  already  referred  to.  By  no 
process  known  to  man  can  it  be  decomposed  into  anything 
simpler  than  copper. 

On  the  other  hand,  take  the  well-known  substance  water,  add 
a  little  acid  to  it  to  make  it  conduct,  and  pass  an  electric  cur- 
rent through  it.  The  water  will  be  decomposed  into  two  simpler 
substances,  both  of  them  gases,  and  they  will  be  set  free  the 
one  at  the  one  pole,  the  other  at  the  other. 

There  is,  therefore,  a  fundamental  difference  between  copper 
and  water  —  the  one  cannot  be  decomposed  into  simpler  sub- 
stances, the  other  can  be  decomposed  into  two  substances,  both 
of  which  differ  fundamentally  in  their  properties  from  the  sub- 
stance water. 

Substances  like  copper  which  cannot  be  decomposed  into 
anything  simpler  are  known  as  elements,  while  those  substances 
which  can  be  decomposed  into  simpler  things  are  known  as 
compounds. 

The  Number  of  Elements  and  Compounds. — While  the  chemi- 
cal compounds  already  known  number  more  than  a  hundred 
thousand,  the  number  of  chemical  elements  which  have  thus 
far  been  discovered  is  only  about  seventy-five.  When  we  con- 


ELEMENTS  OF  INORGANIC  CHEMISTRY 


sider  that  a  compound  is  made  up  of  two  or  more  of  these 
elementary  substances,  the  whole  problem  of  the  composition 
of  substances  is  vastly  simplified.  We  can,  then,  refer  every 
compound  known,  both  in  inanimate  nature  and  in  the  realm 
of  living  matter,  to  a  comparatively  few  elementary  substances. 

Take  the  rocks  that  are  most  familiar  on  the  surface  of  the 
earth;  they  are  made  up  chiefly  of  not  more  than  a  dozen 
elementary  substances. 

Water,  which  covers  such  a  large  portion  of  the  surface  of 
the  earth,  is,  as  we  have  seen,  made  up  of  two  elements. 

The  atmosphere  which  is  so  essential  to  life  is  made  up 
chiefly  of  two  elements. 

If  we  turn  to  living  matter,  we  find  a  very  large  number  of 
chemical  compounds,  and  the  greatest  complexity  represented. 
Cellulose,  starch,  albumen,  are  among  the  most  complex  sub- 
stances known  to  the  chemist;  yet,  an  analysis  of  these 
substances  brings  out  the  surprising  fact  that  they  contain 
scarcely  more  than  a  half-dozen  elements. 

The  number  of  chemical  elements  known  to  us  at  present  is, 
as  already  stated,  about  seventy-five. 

The  Chemical  Elements.  —  Having  learned  what  is  meant  by 
the  term  "  chemical  element,"  we  naturally  ask  which  are  the 
elements  and  what  substances  are  compounds  ?  In  the  follow- 
ing table  the  substances  which  have  been  shown  with  a  reason- 
able degree  of  probability  to  be  elementary,  are  given,  together 
with  the  symbol  which  is  used  for  the  element  in  question.  The 
elements  in  heavy  type  are  either  the  most  abundant  in  nature, 
or  the  most  important  chemically.  Those  in  italics  are  next  in 
order  of  importance,  while  those  printed  in  ordinary  type  are 
either  rare  in  nature  or  are  chemically  of  subordinate  importance. 


Aluminium Al 

Antimony Sb 

Argon A 

Arsenic As 

Barium  Ba 


Bismuth Bi 

Boron B 

Bromine Br 

Cadmium Cd 

Caesium  .     .  Cs 


THE   SCIENCE  OF  CHEMISTRY 


Calcium Ca 

Carbon     .......  C 

Cerium Ce 

Chlorine Cl 

Chromium Cr 

Cobalt Co 

Coliimbiimi  .     .     .     .     .  Cb 

Copper Cu 

Erbium  (?) E 

Fluorine F 

Gadolinium G 

Gallium Ga 

Germanium Ge 

Glucinum Gl 

Gold Au 

Helium He 

Hydrogen H 

Indium In 

Iodine I 

Iridium Ir 

Iron     .......  Fe 

Krypton Kr 

Lanthanum La 

Lead Pb 

Lithium Li 

Magnesium Mg 

Manganese Mn 

Mercury Hg 

Molybdenum     ....  Mo 

Neodymium Nd 

Neon Ne 

Nickel .  Ni 

Nitrogen N 


Osmium Os 

Oxygen O 

Palladium Pd 

Phosphorus P 

Platinum Pt 

Potassium K 

Proseodymium .     .     .     .  Pr 

Ehodium Eh 

Rubidium Eb 

Ruthenium Eu 

Samarium Sm 

Scandium Sc 

Selenium Se 

Silicon Si 

Silver Ag 

Sodium Na 

Strontium Sr 

Sulphur S 

Tantalum Ta 

Tellurium Te 

Thalium Tl 

Thorium Th 

Thulium Tu 

Tin.     .     .     ...     .     .  Sn 

Titanium Ti 

Tungsten W 

Uranium U 

Vanadium V 

Xenon X 

Ytterbium Yb 

Yttrium Y 

Zinc Zn 

Zirconium  Zr 


The  symbol  of  the  element  is  usually  the  first  letter,  or  this 
combined  with  some  other  distinctive  letter  when  several  ele- 


6  ELEMENTS  OF  INORGANIC   CHEMISTRY 

ments  begin  with  the  same  letter.  In  some  cases,  however, 
the  symbol  is  not  taken  from  the  English  name  of  the  element, 
but  from  the  Latin.  Thus,  the  symbol  for  copper  is  Cu,  from 
the  Latin  cuprum  ;  the  symbol  for  iron  Fe,  from  the  Latin 
ferrum  ;  etc. 

Some  elements  occur  in  very  large  quantities  in  the  earth, 
while  others  are  comparatively  rare.  The  following  estimate 
of  the  composition  of  that  part  of  the  earth  which  is  accessible 
to  us  seems  on  the  whole  the  most  reliable :  — 

Oxygen,  percentage  in  the  earth 50.0 

Silicon,  "  "  "       " 25.0 

Aluminium,  «  "  "       " 7.2 

Iron,  «  "  "       « 5.0 

Calcium,  "  «  "       « 3.5 

Magnesium,  "  "  "       " 2.5 

Sodium,  "  "  "       " 2.3 

Potassium,  «  "  "       " 2.2 

Hydrogen,  "  "  "       « 1.0 

Titanium,  "  «  «  "......  0.3 

Carbon,  "  "  "       " 0.2 

The  earth  is  thus  made  up  chiefly  of  nine  elements,  the 
remainder  occurring  in  comparatively  small  quantities. 

Chemical  Combination.  —  Certain  elements  can  combine  with 
certain  other  elements  and  form  compounds.  Two  elements 
may  combine  and  form  a  compound,  or  three,  four,  or  more 
elements  may  combine. 

While  elementary  substances  may  combine  with  one  another 
and  form  compounds,  it  is  not  true  that  any  element  can  com- 
bine with  any  other  element.  We  shall  learn  that  elements  with 
widely  different  properties  generally  combine  most  readily. 

The  science  of  chemistry  consists,  in  part,  of  a  study  of  the 
elements  and  the  compounds  which  these  elements  can  form 
with  one  another.  Certain  generalizations  have  been  reached 
to  which  chemical  reactions  between  substances  conform. 
Some  of  these  will  be  considered  in  the  proper  places. 


CHAPTER  II 

OXYGEN 

Occurrence  in  Nature.  —  Oxygen  is  the  most  abundant  of  all 
the  chemical  elements.  It  forms  88.8  per  cent  of  all  the 
water  on  the  earth,  and  about  23  per  cent  of  the  atmospheric 
air.  It  is  an  important  constituent  of  most  of  the  rocks,  and 
occurs  in  nearly  all  living  matter  whether  vegetable  or  animal. 
It  is  estimated  in  general  that  about  one-half  of  the  earth's 
crust  is  composed  of  the  element  oxygen. 

Preparation  of  Oxygen.  —  Since  oxygen  occurs  in  such  large 
quantities  in  nature,  we  would  think  that  we  should  turn  to 
some  natural  source  for  a  supply  of  this  element.  It  is,  how- 
ever, not  very  easy  to  obtain  pure  oxygen  from  any  natural 
source.  One  of  the  most  convenient  means  of  obtaining 
oxygen  in  the  laboratory  is  by  heating  potassium  chlorate. 
This  compound  contains  about  39  per  cent  of  oxygen,  and 
gives  up  all  of  its  oxygen  when  moderately  heated. 

Another  method  by  which  oxygen  can  be  obtained  is  to  heat 
manganese  dioxide,  a  compound  of  manganese  and  oxygen 
which  is  rich  in  oxygen.  When  this  compound  is  heated  to 
a  high  temperature  it  gives  up  a  part  of  its  oxygen. 

A  better  method  for  preparing  oxygen  is  by  heating  a  mixture 
of  potassium  chlorate  and  manganese  dioxide.  Oxygen  is  evolved 
in  large  quantity  at  a  comparatively  low  temperature,  and,  on 
the  whole,  this  is  the  best  method  for  preparing  oxygen  in  the 
laboratory. 

Oxygen  can  also  be  prepared  by  heating  a  compound  of  mer- 
cury and  oxygen  known  as  mercuric  oxide.  This  compound  is 
decomposed  by  heat  into  its  two  elements,  oxygen  and  mercury. 

7 


8  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Oxygen  can  be  obtained  at  ordinary  temperatures  by  bringing 
the  compound  of  oxygen  and  hydrogen  known  as  hydrogen  dioxide 
in  contact  with  manganese  dioxide  or  lead  dioxide  in  the  presence 
of  an  acid.  Under  these  conditions  one-half  of  the  oxygen 
liberated  conies  from  the  hydrogen  dioxide,  and  the  other  half 
from  the  manganese  dioxide  or  the  lead  dioxide.  This  differs 
from  the  preceding  methods  in  that  the  oxygen  is  obtained  at 
ordinary  temperatures. 

Substances  burn  readily  in  Oxygen.  —  One  of  the  most  char- 
acteristic of  the  chemical  properties  of  oxygen  is  the  readiness 
with  which  substances  burn  in  it.  Substances  which  burn 
comparatively  slowly,  or  will  not  burn  at  all  in  the  air,  often 
burn  with  the  greatest  readiness  in  oxygen  gas,  emitting  very 
bright  light  and  evolving  large  quantities  of  heat. 

When  a  splinter  of  wood  which  has  been  burned  in  part  to  a 
coal,  and  is  still  glowing,  is  plunged  into  a  vessel  filled  with 
oxygen  gas,  it  bursts  again  into  flame. 

When  sulphur  is  ignited  in  the  air,  it  burns  quietly  with  a 
pale  blue  flame.  When  the  burning  sulphur  is  plunged  into 
a  vessel  filled  with  oxygen,  it  bursts  into  violent  combustion, 
evolving  large  amounts  of  heat  and  light. 

A  steel  watch-spring,  or  picture  wire,  dipped  into  molten  sul- 
phur and  ignited,  will  not  continue  to  burn  in  the  air.  If  after 
the  sulphur  is  ignited  the  steel  watch-spring  is  plunged  into 
oxygen  gas,  the  iron  will  burn  with  an  intense  white  light,  and 
a  large  number  of  highly  heated  particles  will  fly  off  from  the 
iron,  producing  quite  a  pyrotechnic  effect. 

While  phosphorus  burns  quietly  in  the  air,  in  pure  oxygen  it 
burns  with  great  violence.  This  will  be  seen  to  be  the  case 
if  a  piece  of  ignited  phosphorus  is  plunged  into  a  vessel  filled 
with  pure  oxygen. 

Explanation  of  the  Above  Results.  —  The  above  results  show 
beyond  question  that  substances  which  burn  slowly  in  the  air, 
or  do  not  burn  at  all,  burn  readily  in  pure  oxygen.  This  natu- 
rally raises  the  question  why  is  this  the  case  ?  The  air,  as  we 
shall  learn,  is  essentially  oxygen  diluted  with  about  four  times 


OXYGEN 


9 


its  volume  of  nitrogen.  The  number  of  oxygen  particles  in  a 
given  volume  of  air  is,  therefore,  much  less  than  in  a  given 
volume  of  pure  oxygen.  The  nitrogen  serves  to  dilute  the 
oxygen.  When  combustion  takes  place  in  pure  oxygen,  the 
heat  that  is  liberated  is  expended  in  raising  the  temperature 
of  the  oxygen  alone,  and  the  rapidity  of  the  combustion  de- 
pends chiefly  upon  the  temperature  of  the  oxygen  gas. 

When  the  oxygen  is  diluted  with  an  inert  gas  like  nitrogen, 
much  of  the  heat  that  is  set  free  during  the  combustion  is  ex- 
pended in  raising  the  temperature  of  the  nitrogen  which  takes 
no  part  in  the  combustion,  and  as  far  as  accelerating  the  com- 
bustion is  concerned  is,  therefore,  lost. 


COMBUSTION 

The  subject  of  combustion,  or  burning,  is  one  that  has 
attracted  the  attention  of  chemists  from  Very  early  times. 
This  would  be  expected,  since  com- 
bustion is  among  the  most  familiar 
of  chemical  phenomena. 

Oxygen  used  up  in  Combustion. — 
That  oxygen  is  actually  used  up  in 
combustion  can  be  readily  shown  as 
follows :  Fill  a  glass  tube  with  air 
and  introduce  a  piece  of  phosphorus 
into  the  tube  as  shown  in  Figure  1. 
The  phosphorus  will  undergo  slow 
combustion,  and  that  the  oxygen  is 
being  used  up  is  shown  by  the  fact 
that  after  a  slight  expansion  due  to 
heating  has  occurred,  the  water  will 
rise  slowly  but  steadily  in  the  tube. 
About  one-fifth  of  the  air  is  thus 
consumed. 

Increase  in  Weight  in  Combustion.  —  If  combustion  consists 
in  the  union  of  oxygen  with  the  substance  burned,  then  the 


FIG.  1. 


10 


ELEMENTS   OF  INORGANIC  CHEMISTRY 


weight  of  the  products  of  combustion  must  be  greater  than 
the  weight  of  the  substance  that  has  been  burned. 

That  the  products  of  combustion  weigh  more  than  the  sub- 
stance before  it  was  burned  can  be  readily  shown  by  the  follow- 
ing experiment.  Two  pieces  of  caudle  of  equal  length  are 
placed,  one  upon  each  pan  of  a  large  balance.  A  lamp  chimney 

is  suspended  from 
each  end  of  the  arm 
of  the  balance.  A 
piece  of  wire  gauze 
which  fits  the  chim- 
ney tightly  is  in- 
troduced into  each 
chimney,  and  some 
coarse  pieces  of 
caustic  soda  added. 
Caustic  soda  is 
now  added  to  the 
lighter  side  until 
the  pointer  stands 
exactly  in  the  mid- 
dle of  the  scale. 
One  of  the  candles  is  now  lighted  and  the  products  of  com- 
bustion, carbon  dioxide  and  water,  are  caught  by  the  caustic 
soda.  After  the  candle  has  burned  for  a  time  this  arm  of  the 
balance  will  begin  to  sink,  showing  that  the  products  of  com- 
bustion of  the  candle  are  heavier  than  the  unburned  candle. 

Combustion  consists  in  the  union  of  oxygen  with  the  substance 
burned.  We  have  just  seen  that  oxygen  is  used  up  in  com- 
bustion. Also  that  the  weight  of  the  products  of  combustion 
is  greater  than  the  weight  of  the  substance  before  it  was 
burned.  If  we  should  weigh  accurately  the  substance  before 
it  was  burned,  and  the  amount  of  oxygen  used  up  in  the  com- 
bustion, we  would  find  that  the  sum  of  these  two  weights  was 
always  exactly  equal  to  the  weight  of  the  products  of  the  com- 
bustion. This  shows  that  combustion  consists  in  the  union  of 


FIG.  2. 


OXYGEN  11 

the  substance  burned  with  oxygen,  and  combustion  is  nothing 
but  oxidation. 

Rapid  and  Slow  Oxidation.  —  Combustion,  as  we  ordinarily 
observe  it,  is  a  comparatively  rapid  process.  The  substance 
burns  up,  as  we  say,  in  a  few  minutes,  and  there  is  usually  a 
large  evolution  of  heat,  and  in  many  cases  a  marked  production 
of  light.  This  is  known  as  rapid  oxidation. 

We  know  oxidation  processes,  however,  which  take  place 
slowly  and  extend  over  long  periods  of  time,  even  years.  Ex- 
amples are  the  oxidation  of  metals,  the  decaying  or  slow 
oxidation  of  wood,  and  the  like. 

When  the  oxidation  proceeds  slowly,  as  in  these  cases,  there 
is  no  apparent  evolution  of  heat  and  no  evolution  of  light. 
The  question  arises,  Are  we  justified  in  concluding  that  there 
is  actually  no  evolution  of  heat  when  slow  oxidation  takes 
place?  We  cannot  detect  any  heat  set  free,  yet  it  might 
be  that  there  is  a  slow  evolution  of  heat,  but  so  slow  that  it 
escapes  before  it  can  be  detected. 

While  we  cannot  prove  directly,  unless  large  masses  of  sub- 
stances are  employed,  that  heat  is  set  free  in  slow  oxidation, 
it  can  be  proved  indirectly  by  methods  which  it  would  lead  us 
too  far  at  present  to  discuss  in  detail.  Exactly  the  same  amount 
of  heat  is  evolved  when  a  given  amount  of  substance  is  oxi- 
dized, whether  the  combustion  takes  place  slowly  or  rapidly. 

Measurement  of  the  Heat  of  Combustion.  —  To  measure  the 
amount  of  heat  set  free  in  any  chemical  reaction,  such  as  com- 
bustion, the  reaction  must  be  carried  out  in  a  vessel  surrounded 
by  a  poor  conductor  of  heat,  so  that  the  loss  in  heat  will  be 
reduced  to  a  minimum.  The  heat  that  is  produced  is  allowed 
to  warm  a  known  weight  of  water,  and  the  temperature  of  the 
water  is  noted  before  and  after  the  experiment.  The  apparatus 
which  is  used  for  measuring  quantity  of  heat  is  known  as  a 
calorimeter. 

Some  unit  must  be  adopted  for  expressing  the  results  of 
calorimetric  measurements.  Whatever  unit  we  select  would 
be  purely  arbitrary.  The  amount  of  heat  required  to  raise  one 


12  ELEMENTS   OF   INORGANIC   CHEMISTRY 

gram  of  water  from  0°  to  1°  C.  is  frequently  taken  as  the 
unit,  and  is  called  the  calorie  and  written  cal 

Heat  of  Formation  and  of  Decomposition.  —  We  have  just 
seen  that  when  two  or  more  substances  unite  and  form  a  third 
substance,  heat  is  evolved.  Further,  a  definite  amount  of  heat 
is  set  free  when  a  given  amount  of  any  substance  is  formed. 
This  amount  is  known  as  the  heat  of  formation  of  the  substance. 

Given  a  substance  already  formed  by  the  union  of  two  or 
more  substances.  A  certain  amount  of  heat  must  be  added  to 
it  to  decompose  it  into  its  elements.  This  is  known  as  the 
heat  of  decomposition  of  the  substance. 

A  very  beautiful  relation  has  been  established  between  the 
heat  of  formation  of  a  substance  and  its  heat  of  decomposition. 
TJie  two  are  equal. 

Certain  Physical  Properties  of  the  Element  Oxygen.  —  Oxygen 
under  ordinary  conditions  is  a  transparent,  colorless,  odorless 
gas.  It  is  slightly  heavier  than  air. 

Oxygen  at  zero  degrees  and  under  what  we  call  normal 
barometric  pressure,  i.e.  760  mm.  of  mercury,  weighs  1.4296 
grams.  Oxygen  is  only  slightly  soluble  in  water,  and  can, 
therefore,  be  collected  in  vessels  over  water. 

The  Pressure  of  Oxygen  varies  with  the  Conditions. — We 
have  referred  to  the  weight  of  a  litre  of  oxygen  under  normal 
conditions  of  temperature  and  pressure.  This  would  imply 
that  the  weight  of  a  litre  of  oxygen  would  change  if  we  changed 
temperature  or  pressure,  and  such  is  the  fact.  If  we  have  a 
litre  of  oxygen  at  any  given  pressure  and  subject  the  gas  to  a 
..greater  pressure,  the  volume  would  be  less  than  a  litre,  and, 
•Consequently,  the  density  of  the  gas  would  be  increased  and 
|he  weight  of  a  given  volume  of  the  gas  increased.  Similarly, 
•diminution  in  pressure  would  cause  increase  in  volume  and, 
consequently,  diminution  in  the  weight  of  a  given  volume  of 
the  gas. 

If  instead  of  varying  the  pressure  we  vary  the  temperature 
to  which  the  oxygen  gas  is  subjected,  we  would  also  produce 
change  in  volume.  If  the  temperature  of  the  gas  is  increased 


OXYGEN  13 

and  the  pressure  kept  constant,  the  volume  of  the  gas  would 
increase.  If,  on  the  other  hand,  the  temperature  of  the  gas  is 
lowered,  the  pressure  being  kept  constant,  the  volume  of  the 
gas  would  be  diminished.  Certain  relations  between  the  press- 
ure and  volume,  and  the  temperature  and  volume  of  not  only 
oxygen  gas,  but  of  gases  in  general,  have  been  established,  but 
these  will  be  considered  later. 

The  Liquefaction  of  Oxygen.  —  Although  oxygen  is  a  gas 
under  atmospheric  pressure  and  at  all  ordinary  temperatures, 
it  does  not  follow  that  it  is  a  gas  at  all  temperatures  and  press- 
ures. If  we  look  into  the  history  of  the  liquefaction  of  gases,  we 
find,  however,  that  oxygen  resisted  for  a  long  time  all  efforts  to 
liquefy  it,  and  was  placed  among  the  so-called  permanent  gases. 

It  is  now  readily  liquefied  and  in  comparatively  large 
quantity.  When  oxygen  gas  is  cooled  below  a  certain  tem- 
perature (— 119°),  and  then  subjected  to  high  pressure,  it  passes 
into  a  liquid.  The  actual  method  of  liquefying  oxygen  is,  how- 
ever, more  complicated.  Liquid  oxygen  i$  light  blue  in  color, 
and  boils  at  — 181°  0.  under  the  pressure  of  the  atmosphere. 
It  is  a  general  rule  that  liquids  boil  under  diminished  press- 
ure at  a  lower  temperature  than  under  normal  atmospheric 
pressure.  By  placing  liquid  oxygen  under  greatly  diminished 
pressure  a  temperature  as  low  as  —  225°  can  be  obtained. 

EXPERIMENTS  WITH  OXYGEN 

Experiment  1.  Preparation  of  Oxygen  from  Potassium  Chlo- 
rate. —  (Glass  retort  holding  100  or  150  cc. ;  rubber  tube ;  glass 
tube ;  pneumatic  trough ;  glass  cylinders ;  potassium  chlorate.) 

Into  a  small  retort  R  (Fig.  3)  introduce  about  3  grams  of  potas- 
sium chlorate.  Connect  the  retort  by  means  of  a  rubber  tube, 
with  a  glass  tube  A,  bent  as  shown  in  the  figure.  The  trough 
T  is  filled  with  water  until  the  end  of  the  tube  A  is  covered 
with  water.  Heat  the  retort  containing  the  potassium  chlo- 
rate gently  by  means  of  a  bunsen  burner  B.  The  potassium 
chlorate  at  first  melts,  and  then  at  a  somewhat  higher  tempera- 
ture a  gas  begins  to  escape.  After  enough  gas  has  been  gen- 
erated to  expel  the  air  from  the  retort  and  tube,  bring  the 


14 


ELEMENTS   OF   INORGANIC   CHEMISTRY 


glass  cylinder  O,  filled  with  water,  over  the  end  of  the  tube  A, 
as  shown  in  the  figure.  When  a  considerable  amount  of  gas 
has  been  liberated,  the  contents  of  the  retort  will  become  solid, 
and  the  escape  of  gas  from  the  tube  A  will  practically  cease. 
If  now  the  temperature  is  raised  considerably  higher,  oxygen 
will  again  be  given  off  and  in  much  larger  quantity  than  in  the 
first  stage  of  the  reaction.  After  a  time  oxygen  will  cease 
to  be  given  off,  and  no  matter  how  high  the  temperature  is 
raised,  it  is  impossible  to  obtain  any  more  oxygen  from  the 
solid  which  remains  in  the  retort.  The  residue  is  potassium 


chloride,  a  compound  containing  only  potassium  and  chlorine, 
and  no  oxygen.  Prove  that  the  gas  is  oxygen  by  means  of  a 
glowing  splinter.  This  is  obtained  by  igniting  a  splinter  in  a 
bunsen  burner  until  the  end  is  burned  to  a  coal.  Then  extin- 
guish the  flame  and  plunge  the  splinter  quickly  into  the 
vessel  containing  the  oxygen  gas.  The  splinter  will  be  re- 
kindled into  flame  by  the  gas.  (See  page  7.) 

Experiment  2.  Preparation  of  Oxygen  from  Manganese  Di- 
oxide.—  (Hard-glass  tube  closed  at  one  end ;  manganese  dioxide.) 

Introduce  a  few  grams  of  manganese  dioxide  into  a  hard-glass 
tube  closed  at  one  end.  Heat  to  a  high  temperature  in  a 
bunsen  burner,  and  show  that  oxygen  is  escaping  by  inserting 
into  the  mouth  of  the  tube  an  ignited  splinter  from  which  the 
flame  has  been  extinguished.  It  will  glow  brightly  in  the 
oxygen  which  escapes  from  the  tube.  (See  page  7.) 


OXYGEN 


15 


Experiment  3.  Preparation  of  Oxygen  from  a  Mixture  of 
Potassium  Chlorate  and  Manganese  Dioxide.  —  (Glass  retort 
holding  200  to  250  cc. ;  rubber  tube;  glass  tube;  pneumatic 
trough;  cylinders;  wide-mouthed  bottles;  bell-jars;  etc.,  man- 
ganese dioxide ;  potassium  chlorate.) 

Mix  25  grains  of  manganese  dioxide  with  25  grams  of  coarsely 
pulverized  potassium  chlorate.  Heat  a  little  of  the  mixture  in 
a  test-tube  to  see  that  there  is  no  explosion.  If  the  manganese 
dioxide  contains  any  organic  matter  mixed  with  it,  as  not 
infrequently  happens,  an  explosion  may  result  when  the  mix- 
ture is  heated.  If  the  oxygen  is  evolved  quietly  from  the 
mixture  when  the  sample  is  heated  and  there  is  no  sign  of  an 
explosion,  the  mixture  is  introduced  into  the  retort.  The 
apparatus  is  arranged  as  in  Figure  3.  The  retort  is  gently 
heated,  when  oxygen  will  begin  to  escape.  Allow  sufficient 
oxygen  to  escape  to  drive  the  air  out  of  the  apparatus ;  then 
fill  a  number  of  cylinders,  bell-jars,  bottles,  etc.,  with  the 
gas.  It  is  also  convenient  to  fill  some  form  of  large  gasometer 
with  oxygen  for  future  use.  When  all  the  oxygen  which  is 
desired  has  been  obtained,  the  tube  is  removed  from  the  trough 
to  prevent  the  entrance  of  the  water  into  the  retort  when  the 
apparatus  is  allowed  to  cool.  The  flame  is  then  removed  from 
the  retort,  when  the  evolution  of  oxygen  quickly  ceases.  (See 
page  7.) 

Experiment  4.  Preparation  of  Oxygen  from  Mercuric  Oxide.  — 
(Glass  tube  closed  at  one  end;  rubber  stopper;  small  glass  tube; 
pneumatic  trough;  glass  cylinder;  mercuric  oxide.) 


FIG.  4. 


16  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Introduce  a  few  grams  of  mercuric  oxide  into  the  glass  tube 
A  (Fig.  4).  This  tube  is  connected  with  the  cylinder  C  by 
means  of  the  glass  tube  B.  The  cylinder  is  filled  with  water 
and  inverted  in  the  pneumatic  trough  D,  also  containing  water. 
When  the  oxide  is  heated  by  means  of  a  bunsen  burner,  oxygen 
is  set  free  and  collects  in  the  cylinder,  while  mercury  condenses 
upon  the  colder  portions  of  the  tube  A.  Introduce  an  ignited 
splinter  into  the  cylinder.  The  splinter  will  be  quickly  re- 
kindled into  flame  by  the  oxygen  gas.  (See  page  7.) 

Experiment  5.  Preparation  of  Oxygen  from  a  Mixture  of 
Manganese  or  Lead  Dioxide,  and  Hydrogen  Dioxide. — -(Test- 
tube;  manganese  or  lead  dioxide;  hydrogen  dioxide;  dilute  sul- 
phuric acid.) 

Introduce  a  few  grams  of  manganese  dioxide  into  a  test- 
tube,  and  a  few  cubic  centimetres  of  hydrogen  dioxide  and 
dilute  sulphuric  acid.  A  copious  evolution  of  oxygen  will 
take  place.  This  can  be  collected  in  a  cylinder  as  in  Experi- 
ment 4,  or  its  presence  demonstrated  by  means  of  an  ignited 
taper.  (See  page  8.) 

Experiment  6.  Combustion  of  Carbon  in  Oxygen.  —  (A 
cylinder  filled  with  oxygen ;  a  piece  of  charcoal ;  preferably  a 
porous  variety.) 

That  carbon  will  burn  in  oxygen  more  readily  than  in  the  air 
has  already  been  seen  by  plunging  a  splinter  from  which  the 
flame  has  been  extinguished  into  a  vessel  containing  oxygen 
gas.  A  piece  of  carbon  wrapped  with  copper  wire,  which 
serves  to  hold  it,  is  heated  in  a  bunsen  burner  to  redness,  and 
quickly  plunged  into  the  vessel  containing  oxygen.  It  glows 
brightly  in  the  oxygen,  showing  far  more  vigorous  combustion 
than  in  the  air.  (See  page  8.) 

Experiment  7.  Combustion  of  Sulphur  in  Oxygen.  —  (A  defla- 
grating spoon;  a  wide-mouthed  bottle  or  bell-jar  holding  2 
or  3  litres  of  oxygen;  a  piece  of  sulphur.) 

Ignite  the  sulphur  in  a  deflagrating  spoon,  and  observe  that 
it  burns  quietly  in  the  air  with  a  pale  blue  flame.  Plunge  the 
burning  sulphur  into  the  vessel  containing  oxygen.  The 
sulphur  burns  in  the  oxygen  far  more  vigorously  than  in  the 
air.  This  is  shown  by  the  character  of  the  flame,  which  is 
now  completely  changed  both  in  magnitude  and  color.  Instead 
of  the  small,  pale  blue  flame,  which  was  produced  when  the 
sulphur  was  burned  in  the  air,  we  have  a  large,  very  bright 
flame  with  more  or  less  scintillation.  Open  the  vessel  and 
carefully  smell  the  contents.  It  has  the  odor  of  a  burning 


OXYGEN 


17 


FIG.  5. 


sulphur  match.,  and  thus  differs  from 
the  original  oxygen.     (See  page  8.) 

Experiment  8.  Combustion  of  Iron 
in  Oxygen.  —  (Steel  watch-spring,  or 
picture  wire;  cotton  thread;  vessel 
holding  3  or  4  litres  of  oxygen;  sul- 
phur.) 

Wind  the  lower  end  of  an  old  steel 
watch-spring  with  cotton  and  dip  it 
into  molten  sulphur.  The  sulphur  ad- 
heres to  the  cotton  in  considerable 
quantity.  Ignite  the  sulphur  and  plunge 
the  spring  into  oxygen.  The  sulphur 
will  burn  vigorously  and  heat  the  steel 
to  the  temperature  at  which  it  will 
begin  to  unite  with  oxygen.  The 
iron  will  then  continue  to  burn  in  the 
oxygen,  emitting  an  intense  light,  and 
white-hot  particles  of  iron  will  fly  off 

in  all  directions  (Fig.  5).  If  a  closed  vessel  is  used  to  contain 
the  oxygen  in  this  experiment,  it  is  advisable  to  cover  the  bot- 
tom with  water,  in  order  that  the  hot  particles  of  iron  falling 
upon  it  may  not  crack  the  vessel.  (See  page  8.) 

\  Experiment  9.    Combustion  of  Phos- 

phorus in  Oxygen.  —  (A  wide-mouthed 
vessel  containing  several  litres  of  oxy- 
gen; a  deflagrating  spoon;  a  small 
piece  of  phosphorus.) 

Introduce  a  very  small  piece  of  phos- 
phorus (3  or  4  mm.  in  diameter)  into 
a  deflagrating  spoon,  ignite  and  intro- 
duce into  the  oxygen.  The  phosphorus 
will  burn  with  intense  light  and  heat 
(Fig.  6).  It  is  well  to  wrap  the  vessel 
containing  the  oxygen  with  a  towel, 
so  that  in  case  it  should  break  pieces 
of  glass  will  be  prevented  from  fly- 
ing. In  this  experiment  the  precaution 
I  JI1  ||[  should  be  taken  to  cut  the  phosphorus 

III      ....        ID  under  water  and  dry  it  between  pieces 

of    filter-paper.      Phosphorus   should 
never  be  allowed  to  come  in  contact 
FIG.  6.  with  the  hand,  since  bad  wounds  easily 


18  ELEMENTS  OF  INORGANIC  CHEMISTRY 

result.  In  the  above  experiment  special  precaution  should  be 
taken  to  use  only  a  very  small  piece  of  phosphorus.  (See 
page  8.) 

Experiment  10.  Oxygen  is  used  up  in  Combustion.  —  (A 
tall,  slender,  bell-jar;  a  deflagrating  spoon;  sulphur;  steel 
watch-spring.) 

This  experiment  can  be  performed  essentially  as  Experiments 
7,  8,  and  9 ;  the  difference  being  that  the  deflagrating  spoon 
or  the  watch-spring  must  fit  tightly  in  a  rubber  stopper,  which 
is  of  such  size  as  to  close  air-tight  the  neck  of  the  bell-jar 
containing  the  oxygen.  The  bell-jar  rests  in  a  vessel  holding 
a  considerable  amount  of  water.  After  the  sulphur  or  iron  is 
ignited  the  stopper  is  forced  tightly  into  the  neck  of  the  bell- 
jar.  As  the  sulphur  or  iron  burns  the  water  rises  in  the 
bell-jar,  showing  that  the  oxygen  is  being  used  up  in  the 
combustion.  (See  page  9.) 


CHAPTER    III 

HYDROGEN 

Occurrence. — Hydrogen  is  apparently  the  most  widely  dis- 
tributed of  all  the  elements.  It  occurs  in  the  sun,  especially 
in  the  prominences  seen  during  solar  eclipses,  in  the  stars,  and 
even  in  the  nebulous  masses  scattered  throughout  the  universe. 

The  greatest  amount  of  hydrogen  on  the  earth,  by  far,  is  in 
water,  whence  the  name  (hydro,  water,  and  gennao,  to  produce). 
All  water  contains  about  11.2  per  cent  of  hydrogen,  and  when 
we  consider  the  amount  of  water  upon  the  earth,  we  get  some 
idea  of  the  amount  of  hydrogen  present  on  our  globe.  It  also 
occurs  in  "most  forms  of  living  matter. 

Preparation  of  the  Element  Hydrogen.  —  To  obtain  the  ele- 
ment hydrogen,  we  would  naturally  turn  to  water  as  the 
largest  source.  Hydrogen  can  be  obtained  from  water  by 
several  means.  When  a  little  acid  is  added  to  water  and  the 
electric  current  passed  through  the  acidified  water,  hydrogen 
gas  is  liberated  at  one  of  the  poles,  and  can  be  easily  collected. 

Hydrogen  can  also  be  obtained  from  water  by  purely  chemi- 
cal means.  When  metallic  sodium  is  brought  in  contact  with 
water  at  ordinary  temperatures,  a  violent  reaction  takes  place, 
and  the  hydrogen  set  free  is  ignited.  When  potassium  is  used 
instead  of  sodium,  a  still  more  violent  reaction  takes  place,  and 
the  hydrogen  set  free  is  ignited. 

Fin  practice  we  seldom  use  any  of  the  above  methods,  since 
we  have  means  of  preparing  hydrogen  on  a  large  scale  which 
are  far  more  convenient  than  any  of  the  above.  When  zinc  is 
treated  with  a  strong  acid,  such  as  hydrochloric  or  sulphuric, 
the  metal  passes  into  solution  and  the  hydrogen  from  the  acid 
escapes. 

19 


20 


ELEMENTS  OF  INORGANIC  CHEMISTRY 


Hydrogen  is  readily  prepared  as  follows:  Some  pieces  of 
zinc  are  introduced  into  a  glass  flask  A,  as  shown  in  the  figure 
(Fig.  7),  and  dilute  hydrochloric  acid  poured  into  the  flask 
through  the  funnel-tube  B,  until  the  end  of  the  tube  dips  be- 
neath the  acid.  Hydrogen  gas  will  be  liberated  and  escape 
through  the  side  tube  C. 

If  it  is  desired  to  prepare  hydrogen  on  a  still  larger  scale,  a 
form  of  apparatus  devised  by  Kipp  (Fig.  8)  is  very  convenient. 


FIG.  7. 


FIG.  8. 


From  this  apparatus  hydrogen  is  obtained  by  simply  turning  a 
stop-cock.  When  no  more  gas  is  desired,  the  stop-cock  is  closed, 
and  the  pressure  of  the  hydrogen  generated  automatically 
drives  the  acid  away  from  the  zinc  and  prevents  the  further 
liberation  of  gas. 

Combination  of  Hydrogen  with  Oxygen.  —  Hydrogen,  a  color- 
less and  odorless  gas,  combines  readily  with  oxygen  at  elevated 
temperatures.  A  mixture  of  hydrogen  and  oxygen  can  be  kept 
for  an  indefinite  time,  provided  the  mixture  is  not  heated.  If 


HYDROGEN  21 

the  temperature  is  raised  sufficiently,  the  two  combine  with  the 
greatest  ease,  producing  a  violent  explosion. 

That  hydrogen  can  be  burned  in  the  presence  of  oxygen 
without  any  explosion  taking  place  can  be  readily  shown.  A 
rubber  tube  is  attached  to  the  end  of  the  small  glass  tube  C 
(Fig.  7),  and  into  the  other  end  of  the  rubber  tube  a  metallic 
tube  with  a  very  fine  opening  is  inserted.  The  small  tube  at 
the  end  of  a  mouth-blowpipe  works  very  well.  The  hydrogen 
is  allowed  to  escape  from  the  apparatus  through  the  metallic 
tube  until  every  trace  of  air  has  been  removed  from  the  appa- 
ratus. The  hydrogen  is  ignited  at  the  end  of  the  metal  tube. 
It  will  burn  with  a  flame  which  is  nearly  colorless,  but  which 
is  intensely  hot,  as  can  be  shown  by  inserting  a  piece  of  metal 
into  the  flame. 

That  water  is  formed  in  this  process  can  be  readily  demon- 
strated as  follows :  A  cold,  dry,  glass  cylinder  is  brought  over 
the  flame  of  burning  hydrogen,  and  held  in  position  for  a  few 
moments.  The  inner  wall  of  the  cylinder  will  quickly  become 
covered  with  moisture,  and  after  a  short  time  drops  of  water 
will  form  on  the  walls  of  the  cylinder  and  fall  from  the 
mouth. 

The  explosive  nature  of  the  mixture  of  hydrogen  and  oxygen 
can  be  readily  demonstrated  by  the  following  experiment:  Two 
volumes  of  hydrogen  gas  are  mixed  with  one  volume  of  oxygen 
gas,  and  some  of  the  mixture  conducted  through  a  solution  of 
soap  until  a  mass  of  soap  bubbles  has  been  formed.  The  solu- 
tion of  soap  should  be  placed  in  a  thick-walled,  porcelain 
evaporating  dish.  The  dish  containing  the  soap  bubbles  is 
set  in  a  protected  place  such  as  under  the  hood,  and  the 
flame  of  a  gas-lighter  brought  carefully  up  to  the  bubbles  filled 
with  the  mixture  of  hydrogen  and  oxygen.  An  explosion  whose 
violence  depends  on  the  size  and  number  of  the  bubbles  present 
will  take  place.  It  is  well,  therefore,  not  to  have  any  great 
amount  of  the  mixed  gases  present  when  the  flame  is  applied. 

This  mixture  of  the  two  gases  containing  two  volumes  of 
hydrogen  to  one  of  oxygen  is  known  as  detonating  gas  or 


22  ELEMENTS   OF   INORGANIC   CHEMISTRY 

electrolytic  gas,  since  it  is  the  same  mixture  which  is  obtained 
when  an  electric  current  is  passed  through  acidulated  water 
and  the  gases  liberated  at  the  two  poles  allowed  to  mix.  / 

Relations  by  Volume  in  which  Hydrogen  and  Oxygen  Com- 
bine.—  It  was  discovered  early  in  the  nineteenth  century  that 
hydrogen  and  oxygen  combine  in  simple  volume  relations.  No 
matter  in  what  proportions  the  gases  hydrogen  and  oxygen  are 
mixed,  for  every  volume  of  oxygen  that  disappears  when 
combination  takes  place  two  volumes  of  hydrogen  disappear. 
The  ratio  of  the  volumes  which  combine  is,  therefore,  one  to 
two. 

The  further  question  that  remains  is  what  relation  exists 
between  the  volumes  of  the  gases  which  combine  and  the 
volume  of  the  water-vapor  formed  ? 

Two  volumes  of  hydrogen  gas  combine  with  one  volume  of 
oxygen  gas  and  form  two  volumes  of  water-vapor.  Three 
volumes  of  the  constituent  gases  have  disappeared,  and  two 
volumes  of  the  product  have  been  formed.  There  has  been  a 
contraction  in  volume  of  one-third. 

We  shall  learn  from  a  study  of  other  cases  that  this  is  a 
general  relation.  Gases  combine  in  simple  volume  relations, 
and  there  is  a  simple  relation  between  the  volumes  of  the  gases 
which  enter  into  combination  and  the  volume  of  the  product 
formed. 

Heat  Energy  produced  when  Oxygen  and  Hydrogen  Combine. 
—  That  there  is  a  large  amount  of  heat  energy  produced  when 
oxygen  combines  with  hydrogen  is  shown  by  the  fact  that  the 
vessel  which  contains  the  gases  becomes  appreciably  heated. 
It  has  been  utilized  as  a  source  of  very  high  temperature  in 
a  form  of  lamp  which  we  shall  now  describe. 

The  Oxyhydrogen  Blowpipe.  —  The  oxyhydrogen  blowpipe  is 
a  form  of  apparatus  in  which  hydrogen  is  so  burned  in  oxygen 
as  to  concentrate  the  heat  in  a  small  space.  The  apparatus  is 
represented  in  Figure  9.  The  hydrogen  enters  through  the  side- 
tube  H9  and  is  lit  at  E.  Oxygen  enters  through  the  tube  0 
and  does  not  mix  with  the  hydrogen  until  the  flame  is  reached. 


HYDROGEN 


23 


The  flame  of  the  oxyhydrogen  blowpipe  gives  very  little 
light,  but  is  intensely  hot.  It  will  give  some  idea  of  the 
temperature  of  the  flame  to  state  that  platinum  can  be  easily 
melted  in  it. 

While  the  flame  of  the  oxyhydrogen  blowpipe  is  itself  only 
slightly  luminous,  an  intense  light  can  be  produced  by  allowing 
it  to  fall  upon  certain  substances  which  can  be  heated  to  a 
high  temperature  without  fusion.  Such  a  substance  is  ordi- 
nary lime.  When  the  oxyhydrogen  flame  is  allowed  to  fall 
upon  a  cylinder  of  lime,  an  intense  white  light  is  produced. 


FIG.  9. 

This  is  the  Drummond  light.  The  light  is  so  intense  that  it 
can  be  used  where  high  illumination  is  required,  as  in  project- 
ing lanterns  and  the  like. 

The  Reducing  Power  of  Hydrogen.  —  The  tendency  of 
hydrogen  to  combine  with  oxygen  manifests  itself,  not  only 
when  the  oxygen  is  in  the  free  state,  but  even  when  it  is  com- 
bined with  other  elements. 

Take  the  oxide  of  zinc.  When  hydrogen  is  passed  over  this 
substance  at  an  elevated  temperature,  it  combines  with  the 
oxygen  and  leaves  the  zinc  reduced  to  the  elementary  con- 
dition. 

The  removal  of  oxygen  from  a  compound  is  known  as  reduc- 


24  ELEMENTS  OF  INORGANIC   CHEMISTRY 

tion,  and  the   substance  which  can  remove  the  oxygen  as  a 
reducing  agent. 

Nascent  Hydrogen. — When  hydrogen  is  first  liberated  by 
the  action  of  an  acid  on  a  metal,  it  has  very  different  proper- 
ties from  those  which  it  possesses  after  it  has  once  been 
formed.  While  hydrogen  gas,  as  we  ordinarily  know  it,  must 
be  heated  to  an  elevated  temperature  before  it  will  reduce  the 
oxides  of  most  metals,  hydrogen  which  is  just  being  formed 
will  reduce  many  such  substances  even  at  ordinary  tempera- 
tures. Many  other  reactions  which  hydrogen  gas  will  either 
not  effect  at  all,  or  effect  only  at  elevated  temperatures,  will  be 
produced  readily  at  ordinary  temperatures  by  hydrogen  which 
is  just  being  formed. 

Hydrogen  which  is  just  being  formed  has  acquired  a  specific 
name  to  distinguish  it  from  hydrogen  which  has  been  formed 
for  an  appreciable  time.  It  is  known  as  nascent  hydrogen. 
This  condition  of  the  nascent  state  we  shall  learn  is  not  pecul- 
iar to  hydrogen,  but  is  possessed  by  other  elements  as  well. 

Certain  Physical  Properties  of  the  Element  Hydrogen.— 
Hydrogen  is  a  transparent,  colorless  gas,  without  taste  or 
odor.  It  is  the  lightest  of  all  known  substances,  being  nearly 
sixteen  times  lighter  than  oxygen.  One  litre  of  hydrogen  at 
normal  temperature  and  pressure  weighs  only  0.08995  gram. 
The  relative  lightness  or  small  density  of  hydrogen  can  be 
shown  in  a  number  of  ways. 

If  a  small  balloon  or  light  sack  of  any  kind  which  will  hold 
a  gas  is  filled  with  hydrogen,  the  mouth  tied,  and  the  balloon 
set  free,  it  will  rise  rapidly  in  the  air,  showing  that  hydrogen 
is  considerably  lighter  than  air.  This  is  made  use  of  on  a 
large  scale  by  aeronauts  for  ascending  to  considerable  heights 
in  the  atmosphere.  A  large  silk  balloon  is  filled  with  hydro- 
gen, and  it  will  not  only  rise  in  the  atmosphere  but  will  carry 
considerable  weight  with  it.  When  it  is  desired  to  descend, 
the  hydrogen  is  allowed  to  escape  through  a  valve  into  the  air. 

A  still  more  striking  illustration  of  the  small  density  of 
hydrogen  is  shown  by  an  experiment  based  upon  the  rate  at 


HYDROGEN 


25 


which  hydrogen  gas  diffuses.     There  is  a  well-known  law  con- 
necting the  rates  at  which  gases  diffuse  with  their  densities. 

(rases  diffuse  with  velocities  which  are  inversely  proportional 
to  the  square  roots  of  their  densities. 

The  lighter  the  gas,  the  more  rapidly,  then,  will  it  diffuse. 
That  hydrogen  diffuses  rapidly  can  be  shown  by  the  following 
experiment  (Fig.  10) :  A  hollow  porous  cup  C  is  fastened  to 
a  glass  tube  R,  which  ex- 
tends into  the  flask  F, 
passing  through  a  stoppe^ 
which  tightly  closes  the 
mouth  of  the  flask.  A  sec- 
ond glass  tube  T,  drawn 
out  to  a  fine  opening, 
passes  through  a  second 
stopper  and  dips  beneath 
the  water  in  the  flask.  A 
large  bell-jar  F  is  now  filled 
with  hydrogen  and  placed 
over  the  porous,  porcelain 
cup.  Hydrogen  diffuses 
rapidly  in  through  the 
cup,  due  to  the  small  den- 
sity of  the  gas,  produces 
a  pressure  inside  the  ap- 
paratus, and  this  forces 
the  water  up  into  the  glass 
tube  Tj  and  out  through  the  small  opening.  In  this  way  quite 
a  fountain  can  be  produced. 

Hydrogen  is  only  slightly  soluble  in  water,  100  volumes  of 
water  at  15°  dissolving  only  1.9  volumes' of  hydrogen. 

The  Liquefaction  of  Hydrogen.  —  Hydrogen  like  oxygen  was 
one  of  the  few  gases  which  resisted  liquefaction  until  quite 
recently.  When  liquid  oxygen  is  allowed  to  evaporate  under 
low  pressure,  a  temperature  of  —  210°  to  —  225°  can  be  pro- 
duced. If  hydrogen  under  a  pressure  of  several  hundred  atmos- 


FIG.  10. 


26  ELEMENTS   OF  INORGANIC   CHEMISTRY 

pheres  is  cooled  by  liquid  oxygen  to  —  200°  or  —  225°  and  then 
is  suddenly  allowed  to  expand,  it  will  in  expanding  cool  itself 
to  its  point  of  liquefaction. 

The  liquefaction  of  hydrogen  in  appreciable  quantities  we 
owe  almost  entirely  to  Dewar.  He  has  shown  that  its  boiling- 
point  is  —  252°.  It  is,  however,  possible  to  reach  a  still  lower 
temperature  by  a  method  which  has  now  become  familiar  to 
us.  By  allowing  liquid  hydrogen  to  boil  under  greatly  dimin- 
ished pressure,  still  further  cooling  is  produced,  and  a  tem- 
perature as  low  as  —  258°  has  been  realized.  Under  these 
conditions  the  hydrogen  solidified. 

Properties  of  Liquid  Hydrogen.  —  Liquid  hydrogen  is  color- 
less and  transparent,  and  has  small  viscosity.  The  supposed 
blue  color  of  liquid  hydrogen  is  due  to  impurities.  It  has  a 
density  of  0.07,  water  being  unity.  By  contact  with  liquid 
hydrogen,  oxygen  (and  as  we  shall  learn  also  air)  is  converted 
first  into  a  liquid  and  then  into  a  solid,  or  is  frozen,  as  we  say. 

A  beautiful  and  thrilling  experiment  has  been  performed  by 
Dewar,  who  has  liquefied  hydrogen  by  the  litre.  Liquid  hydro- 
gen was  poured  into  a  test-tube  and  the  tube  exposed  to  the  air. 
Liquid  air  soon  began  to  stream  off  the  test-tube,  and  finally  the 
tube  became  covered  with  frozen  air.  The  remarkable  charac- 
ter of  this  experiment  is  evident  to  any  one. 

EXPERIMENTS  WITH  HYDROGEN 

Experiment  11.  Preparation  of  Hydrogen  by  the  Action  of 
Hydrochloric  or  Sulphuric  Acid  on  Zinc.  —  (Glass  flask  of  400 
to  500  cc.  capacity ;  thistle  tube ;  glass  tubing ;  cylinders  for  col- 
lecting gas ;  granulated  zinc  ;  hydrochloric  or  sulphuric  acid.) 

Arrange  the  apparatus  as  shown  in  Figure  7.  Into  the  flask 
A  introduce  some  granulated  zinc,  and  through  the  tube  B  pour 
hydrochloric  acid  or  sulphuric  acid.  If  hydrochloric  acid  is 
used,  a  dilution  of  one  acid  to  two  water  is  convenient.  The 
sulphuric  acid  to  be  used  is  prepared  by  adding  concentrated 
sulphuric  acid  slowly  with  constant  stirring  to  six  times  its 
volume  of  water. 

If  the  zinc  is  fairly  impure,  as  is  ordinarily  the  c$se?  the 


I 


HYDROGEN  27 

rapid  evolution  of  hydrogen  begins  at  once.  Allow  the  hydro- 
gen to  escape  until  every  trace  of  air  has  been  removed  from  the 
apparatus.  Fill  a  test-tube  with  the  gas  and  ignite ;  an  explosion 
shows  that  the  vessel  still  contains  air.  This  is  of  fundamental 
importance,  since  if  there  is  any  oxygen  mixed  with  the  hydro- 
gen, dangerous  explosions  will  result  in  subsequent  experiments. 

Granulated  zinc  is  used  on  account  of  the  comparatively  large 
surface  exposed  to  the  action  of  the  acid.  If  purer  zinc  is  em- 
ployed and  the  action  with  acids  is  slow,  add  a  few  drops  of 
platinic  chloride  to  the  zinc,  when  a  rapid  evolution  of  hydro- 
gen will  take  place. 

The  hydrogen  can  be  readily  dried  by  passing  it  through  a 
cylinder  containing  concentrated  sulphuric  acid.  If  compar- 
atively pure  hydrogen  is  desired,  it  can  be  passed  first  through 
a  solution  of  potassium  permanganate  to  which  a  little  sul- 
phuric acid  has  been  added,  and  then  through  concentrated 
sulphuric  acid  to  remove  the  water. 

Fill  a  number  of  vessels  with  hydrogen  by  displacement  of 
water,  as  with  oxygen.  In  working  with  hydrogen  the  Kipp 
apparatus  (Fig.  8)  is  very  convenient.  The  middle  bulb  is 
partly  filled  with  granulated  zinc,  the  lower  bulb  with  dilute 
sulphuric  acid,  which  extends  also  into  the  upper  bulb.  When 
the  stop-cock  is  open,  the  sulphuric  acid  comes  in  contact  with 
the  zinc  and  hydrogen  is  evolved.  When  the  stop-cock  is 
closed,  the  hydrogen  generated  forces  the  sulphuric  acid  off  of 
the  zinc  into  the  lower  and  upper  bulbs,  and  the  generation  of 
hydrogen  ceases.  When  the  gas  is  wanted,  it  is  only  necessary 
to  open  the  stop-cock  and  obtain  any  desired  quantity.  (See 
page  20.) 

Experiment  12.  Combustion  of  Hydrogen.  —  (Kipp  appara- 
tus for  generating  hydrogen ;  small  tip  of  a  blowpipe ;  platinum 
wire.) 

Connect  a  Kipp  apparatus  for  generating  hydrogen,  or 
the  apparatus  described  in  Experiment  11,  with  the  small  tip 
of  a  mouth-blowpipe,  by  means  of  a  rubber  tube  fastened  very 
tightly  to  the  tip  to  prevent  any  air  from  mixing  with  the 
hydrogen.  Allow  the  hydrogen  to  flow  for  a  time  and  then 
ignite  the  jet.  The  hydrogen  will  burn  with  a  nearly  colorless 
flame,  which,  however,  is  intensely  hot,  as  can  be  shown  by 
holding  in  the  flame  a  piece  of  platinum  wire. 

If  hydrogen  is  burned  from  a  glass  tube,  the  flame  will  be 
colored  yellow,  due  to  the  presence  of  sodium  in  the  glass. 
(See  page  21.) 


28 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


Experiment  13.  Reducing  Power  of  Hydrogen.  —  (Hard- 
glass  tube  closed  at  both  ends  by  one-hole  rubber  stoppers ; 
Kipp  apparatus  for  generating  hydrogen ;  glass  cylinder ;  copper 
oxide ;  sulphuric  acid.) 

Arrange  the  apparatus  as  shown  in  Figure  11.  The  copper 
oxide  is  introduced  into  the  glass  tube  at  (7,  and  the  tube  con- 
nected with  the  drying  cylinder  T,  which,  in  turn,  is  connected 
with  the  hydrogen  generator  by  means  of  a  rubber  tube.  The 
drying  cylinder  is  partly  filled  with  concen- 
trated sulphuric  acid.  The  hydrogen  is  ad- 
mitted through  the  glass  tube  B.  When  the 
hydrogen  has  been  passed  through  the  tube 
until  every  trace  of  air  has  been  removed,  the 
glass  tube  is  heated  at  C  by  means  of  a  bunsen 


FIG.  11. 

burner.  The  copper  oxide  is  gradually  reduced  by  the  hydro- 
gen, yielding  red,  metallic  copper.  Water  collects  in  the  end 
of  the  glass  tube  beyond  the  flame.  (See  page  23.) 

Experiment  14.  Water  formed  when  Hydrogen  is  burned 
in  the  Air.  —  (Hydrogen  generator ;  U-tube  filled  with  calcium 
chloride ;  bent  glass  tube ;  bell-jar  -J-  litre.) 

Arrange  the  apparatus  as  shown  in  Figure  12.  The  hydrogen 
is  passed  through  the  calcium  chloride  to  remove  the  moisture, 
and  then  ignited  at  the  end  of  the  tip.  The  bell-jar  is  placed 
over  the  flame  so  that  the  flame  does  not  touch  the  walls.  As 


HYDROGEN 


29 


FIG.  12. 

the  hydrogen  burns,  water  collects  upon 
the  walls  of  the  bell-jar,  and  drops  from 
the  lower  edge.  (See  page  21.) 

Experiment  15.  Hydrogen  Lighter 
than  Air.  —  (Two  glass  cylinders,  hold- 
ing about  1  litre  each.) 

Fill  one  cylinder  with  hydrogen  and 
cover  the  mouth  with  a  ground-glass 
plate.  Bring  the  mouth  of  this  cylinder 
beneath  the  mouth  of  a  similar  cylinder 
filled  only  with  air,  and  placed  with  its 
mouth  downwards  (Fig.  13).  Hold  the 
cylinders  in  this  position  for  a  few 
seconds,  and  then  apply  a  lighted  match 
to  the  mouths  of  the  two  cylinders.  It 
will  be  found  that  the  hydrogen  has 
passed  from  the  lower  cylinder  into  the 
upper,  —  has  been  poured  upwards, — 
displacing  the  air  in  the  upper  cylinder. 
(See  page  24.) 

Experiment  16.  Hydrogen  Lighter 
than  Air.  —  (Chemical  balance  ;  large 
beaker ;  vessel  holding  from  2  to  3  litres 
of  hydrogen.) 


FIG.  13. 


30 


ELEMENTS   OF  INORGANIC   CHEMISTRY 


Use  tlie  same  chemical  balance  as  in  Figure  2.  Re- 
move both  lamp  chimneys,  replacing  one  by  a  scale  pan,  and  the 
other  by  a  large  beaker  which  is  suspended  from  the  arm  of 
the  balance  in  an  inverted  position.  Add  weights  to  the  scale 
pan  until  the  pointer  stands  at  zero.  Then  allow  the  hydro- 
gen to  flow  from  the 
containing  vessel  up- 
wards into  the  beaker. 
This  arm  of  the  bal- 
ance will  rise,  show- 
ing that  the  hydrogen 
is  lighter  than  the  air 
which  it  displaced. 

This  arm  will  be- 
gin to  fall  again  after 
a  short  time,  due  to 
the  rapid  rate  of  dif- 
fusion of  the  hydrogen 
out  into  the  surround- 
ing air.  (See  page 
24.) 

Experiment  17.  Hy- 
drogen Soap  Bubbles 
rise  in  the  Air.  —  (A 
saturated  solution  of 
soap  containing  about 
one-half  its  volume 
of  glycerine ;  Kipp 
apparatus  for  gener- 
ating hydrogen ;  a 
thistle  tube  or  a  smok- 
er's pipe.) 

The  small  end  of 
the  thistle  tube  or 
pipe  is  introduced  in- 
to the  soap  solution, 
and  the  excess  of  liquid  shaken  off.  The  hydrogen  is  care- 
fully admitted  from  the  Kipp  apparatus,  when  the  soap  bubble 
filled  with  hydrogen  gas  gradually  forms.  By  jarring  the  tube 
or  pipe  the  bubble  is  set  free  and  rises  rapidly  in  the  air. 
By  arranging  a  gas-jet  as  shown  in  Figure  14,  and  inverting  a 
large  funnel  above  it,  the  hydrogen  soap  bubble  is  drawn  into 
the  flame  and  explodes.  (See  page  24.) 


O 


FIG.  14. 


HYDROGEN  31 

Experiment  18.  Diffusion  of  Hydrogen.  —  (Unglazed  porce- 
lain cup  ;  glass  tube ;  glass  bottle ;  bell-jar  filled  with  hydrogen.) 

Arrange  an  apparatus  as  shown  in  Figure  10.  The  tube  It  just 
passes  through  the  rubber  stopper.  The  tube  T  drawn  out  to 
a  fine  point  dips  beneath  the  surface  of  the  liquid.  Bring  the 
vessel  F,  filled  with  hydrogen,  over  the  porous  cup.  Hydrogen 
diffuses  rapidly  through  the  cup,  and  produces  such  a  pressure 
on  the  water  in  the  vessel  that  a  spray  of  water  is  thrown  out 
of  the  fine  opening  of  the  tube  T.  (See  page  25.) 


CHAPTER  IV 

WATER 

Occurrence  of  Water.  —  Water,  on  account  of  its  very  wide 
distribution  over  the  surface  of  the  earth,  is  probably  the  best 
known  chemical  compound.  In  the  free  condition  it  covers 
about  three-fourths  of  the  surface  of  the  earth.  Further,  it  is 
widely  distributed  through  the  rocks  over  the  surface  of  the 
earth,  each  cubic  metre  of  rock  containing  on  the  average  about 
one  litre  of  water.  It  exists  in  large  quantities  in  the  atmos- 
phere, in  the  form  of  water-vapor.  It  also  exists  in  combina- 
tion with  a  large  number  of  substances  as  water  of  crystalli- 
zation, or  water  of  hydration.  Its  presence  is  not  limited  to 
inorganic  or  inanimate  nature.  It  forms  an  essential  part  of  all 
living  matter.  The  main  constituent  of  living  matter  as  far  as 
mass  is  concerned  is  water.  The  human  body  is  more  than 
two-thirds  water,  and  the  animal  and  vegetable  food  which  we 
eat  contains  scarcely  less  water  in  proportion  to  solid  matter. 

Water  as  it  occurs  in  Nature  is  Impure.  —  It  is  safe  to  say 
that  all  natural  water  contains  impurities.  This  does  not  refer 
to  impurities  which  are  thrown  into  water  artificially,  as  by 
the  drainage  of  human  habitations,  but  to  impurities  which  we 
may  call  natural.  The  water  of  the  sea  is  very  impure  because 
of  matter  dissolved  from  the  soil  and  rocks  by  the  waters  be- 
fore they  reach  the  sea,  and  after  they  have  been  poured  into 
it.  The  waters  of  small  streams  and  rivers  are  impure  for  the 
same  reason.  If  water  has  come  in  contact  with  soil  contain- 
ing a  large  amount  of  limestone,  and  especially  if  there  is  much 
organic  matter  in  the  soil,  it  will  dissolve  large  quantities  of 
the  limestone  and  is  then  what  we  call  hard  water. 

32 


WATER  33 

If,  on  the  other  hand,  the  water  has  fallen  upon  a  region 
which  contains  mainly  sandstone  or  other  difficultly  soluble 
rocks,  but  little  of  the  solid  matter  will  dissolve,  and  we  have 
then  comparatively  pure  water.  This  is  the  reason  why  water 
from  mountains  composed  of  granite  or  sandstone  is  relatively 
pure. 

Rain-water  would  undoubtedly  be  fairly  pure  were  it  not 
contaminated  while  in  the  atmosphere.  However,  while  it 
exists  in  the  atmosphere  in  the  form  of  vapor  it  takes  up  many 
kinds  of  impurities,  and  especially  after  it  is  formed  into  drops 
and  falls  through  the  atmosphere,  foreign  matter  is  dissolved 
by  it. 

From  the  above  it  is  then  obvious  that  all  natural  waters 
contain  impurities,  but  that  the  amount  of  impurity  varies 
greatly  from  one  sample  of  water  to  another. 

Mineral  Waters.  —  In  certain  localities  minerals  exist  which 
are  more  or  less  soluble  in  water.  Eain-water  or  water  from 
other  sources  dissolves  these  substances  and  holds  them  in 
solution.  Such  waters  are  known  in  general  as  mineral  ivaters, 
the  nature  of  the  water  depending  upon  the  nature  of  the 
mineral  in  solution. 

Purification  of  Water.  —  Water  is  usually  rendered  impure 
in  the  way  described  above,  by  carrying  with  it  in  solution 
dissolved  substances.  It  may,  however,  be  rendered  impure  by 
matter  which  is  not  in  solution,  but  simply  in  a  state  of  me- 
chanical suspension.  This  latter  condition  is  illustrated  by 
small  streams  after  a  heavy  rain.  The  finely  divided  soil  is 
carried  along  with  the  water  in  a  state  of  fine  suspension,  and 
we  have  muddy  water. 

When  the  impurity  is  in  a  state  of  mechanical  suspension 
and  is  not  in  solution,  it  can  be  removed  by  filtration.  Filtra- 
tion consists  in  passing  water  through  a  substance  with  very 
fine  openings  or  pores,  so  fine,  indeed,  that  the  particles  of 
water  can  pass,  but  not  the  particles  held  in  mechanical  sus- 
pension. This  is  effected  on  a  small  scale  in  the  laboratory  by 
means  of  certain  varieties  of  paper  known  as  "filter-paper" 


34 


ELEMENTS   OF  INORGANIC   CHEMISTRY 


If  the  impurity  in  the  water  is  in  solution,  it  is  obvious  that 
we  cannot  separate  it  by  any  mechanical  process  such  as  filtra- 
tion. Some  other  principle  must  be  utilized.  When  water 
containing  non-volatile  impurities  is  boiled,  the  vapor  which 
escapes  is  practically  pure.  If  this  vapor  is  condensed  again, 
we  have  practically  pure  water.  This  process  of  converting  a 
liquid  into  vapor  and  reconderising  the  vapor  is  known  as  dis- 
tillation, and  the  apparatus  in  which  a  distillation  is  carried 
on  as  a  still. 


Fia.  15. 


The  form  of  still  which  is  used  when  only  a  small  amount 
of  liquid  is  to  be  distilled  is  shown  in  Figure  15.  Into  the  glass 
flask  F  the  liquid  to  be  distilled  is  introduced.  This  is  heated 
and  converted  into  vapor  by  a  burner  placed  beneath  the  flask. 
C  is  the  condenser,  consisting  of  a  small  inner  glass  tube  sur- 
rounded by  a  much  larger  glass-jacket.  Cold  water  is  passed 
into  the  jacket  at  a,  and  out  at  b.  The  vapor  in  the  inner  tube 
is  condensed  to  a  liquid  as  it  passes  through  the  condenser,  and 
flows  into  the  receiver  E. 

If  it  is  desired  to  distil  a  liquid  on  a  large  scale,  the  form  of 


WATER  £5 

the  apparatus  is  greatly  modified,  but  the  principle  is  exactly 
the  same  as  in  the  apparatus  described  above. 

Water  not  an  Element,  but  a  Compound.  —  Water  is  the  first 
substance  that  we  have  thus  far  studied  which  is  not  an  ele- 
ment, but  is  composed  of  more  than  one  element. 

That  water  is  not  an  element  is  obvious  from  our  studies  of 
oxygen  and  hydrogen.  We  have  seen  that  by  electrolysis  both 
oxygen  and  hydrogen  can  be  obtained  from  water ;  and  an  ele- 
ment, by  definition,  is  a  substance  which  cannot  be  decomposed 
into  any  other  substances. 

Composition  of  Water.  —  We  have  seen  that  oxygen  and 
hydrogen  can  be  obtained  from  water,  but  this  does  not  show 
that  water  contains  only  these  two  elements.  To  answer  this 
question  two  general  methods  are  available.  First,  decompose 
water,  and  see  whether  anything  but  hydrogen  and  oxygen  is 
obtained.  Second,  cause  oxygen  and  hydrogen  to  combine,  and 
see  whether  water  is  formed. 

The  most  convenient  means  of  decomposing  water  is  the 
electric  current.  When  a  little  acid  is  added  to  water  to 
diminish  its  resistance  to  the  flow  of  the  current,  and  an 
electric  current  is  passed  through  it,  it  is  decomposed.  This 
process  of  effecting  decompositions  by  means  of  the  current  is 
known  as  electrolysis. 

Electrolysis  of  Water.  —  The  only  products  obtained  by  the 
electrolysis  of  water  are  the  two  gases,  oxygen  and  hydrogen ; 
oxygen  being  set  free  at  the  anode  and  hydrogen  at  the  cathode. 
That  these  are  oxygen  and  hydrogen,  respectively,  can  be 
shown  by  the  fact  that  the  former  will  ignite  a  match  which 
has  just  been  extinguished,  and  the  latter  will  burn  with  the 
characteristic  hydrogen  flame. 

If  we  wish  to  know  the  relative  volumes  of  the  two  gases  set 
free  from  water,  we  must  collect  and  measure  them. 

A  convenient  form  of  apparatus  for  effecting  the  electrolysis 
of  water  and  collecting  the  gases  set  free  is  the  following :  — 

Into  the  two  arms  A  and  B  (Fig.  16)  of  the  U-tube  are  in- 
serted two  platinum  electrodes.  These  tubes  are  completely 


36 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


filled  with,  acidulated  water  by  filling  the  reservoir  R  to  the 
desired  height,  and  opening  the  two  stop-cocks  at  the  ends  of 
A  and  B.  The  current  is  passed  into  the  solution  through  the 
electrode  in  B  and  out  through,  the  electrode  in  A.  The  stop- 
cocks are  closed  be- 
fore the  current  is 
passed,  and  oxy- 
gen collects  in  B, 
and  hydrogen  in  A. 
When  the  current 
has  been  flowing 
for  a  short  time  it 
will  be  observed 
that  the  gas  is  col- 
lecting in  A  faster 
than  in  B.  The 
tubes  A  and  B  are 
graduated  so  that 
at  any  moment  the 
amounts  of  gases 
set  free  can  be 
read  off  at  once. 
After  an  appreci- 
able amount  of  gas 
has  collected  in  B, 
interrupt  the  cur- 
rent and  read  the 
volumes  of  the 
gases  in  the  two 
tubes.  It  will  be 
found  that  there  is 
just  twice  the  volume  of  gas  in  A  that  there  is  in  B.  Close  the 
circuit,  and  allow  the  electric  current  to  flow  until  a  consider- 
ably larger  volume  of  the  two  gases  has  been  set  free.  Inter- 
rupt the  current  again,  and  measure  the  volumes  of  the  two 
gases.  It  will  be  found  that  the  volume  of  the  hydrogen  is 
again  exactly  double  that  of  the  oxygen. 


FIG.  16. 


WATER  37 

No  matter  how  long  the  current  is  allowed  to  flow,  nor  how 
much  water  is  decomposed,  we  would  always  find  that  the 
volume  of  the  hydrogen  set  free  was  exactly  double  that  of 
the  oxygen.  From  the  decomposition  of  water,  or  by  the 
analytical  method,  we  are  therefore  led  to  the  conclusion  that 
water  is  made  up  by  the  union  of  two  volumes  of  hydrogen 
with  one  volume  of  oxygen.  To  determine  the  composition  of 
water,  however,  we  are  not  dependent  solely  upon  the  analytical 
method.  We  can  use  also  the  synthetical. 

Synthesis  of  Water.  —  If  water  is  composed  of  two  volumes 
of  hydrogen  to  one  volume  of  oxygen,  then,  when  we  mix  two 
volumes  of  hydrogen  with  one  volume  of  oxygen  and  pass  an 
electric  spark  through  the  mixture  which  causes  the  gases  to 
combine,  all  the  hydrogen  should  combine  with  all  the  oxygen 
and  form  water.  This  is  exactly  what  takes  place.  Whenever 
two  volumes  of  hydrogen  are  mixed  with  one  volume  of  oxygen 
and  the  gases  made  to  combine  by  means  of  an  electric  spark, 
or  by  rise  in  temperature,  all  the  hydrogen  and  all  the  oxygen 
are  used  up  and  water  is  formed.  If  more  than  two  volumes 
of  hydrogen  are  used,  all  the  oxygen  will  be  used  up  and  the 
excess  of  hydrogen  will  remain  uncombined.  If  less  than  two 
volumes  of  hydrogen  are  used,  all  the  hydrogen  will  be  used 
up  and  the  excess  of  oxygen  will  remain. 

The  results  of  synthesis  confirm  those  of  analysis ;  viz.  that 
water  is  formed  by  the  union  of  two  volumes  of  hydrogen  with 
one  volume  of  oxygen. 

Chemical  Behavior  of  Water.  —  All  things  considered,  water 
is  probably  the  most  important  chemical  compound  known. 
It  is  formed,  as  we  have  seen,  by  the  union  of  hydrogen  and 
oxygen.  Few  compounds  known  are  more  stable  than  water. 
If  we  try  to  decompose  it  into  its  elements,  we  will  appreciate 
what  this  means.  It  is  true,  as  we  have  seen,  that  it  can  be 
decomposed  into  its  elements  by  means  of  the  electric  current. 

If  we  try  to  decompose  water  into  its  elements  by  heat, 
enormous  temperatures  are  required.  In  order  to  effect  even 
slight  decomposition  a  temperature  of  1000°  or  higher  is  neces- 


38  ELEMENTS   OF  INORGANIC   CHEMISTRY 

sary,  and  a  considerable  amount  of  decomposition  is  effected 
only  when  temperatures  between  2000°  and  3000°  are  employed. 

Stress  is  laid  upon  these  facts  for  the  purpose  of  illustrating 
a  general  principle.  When  a  chemical  compound  is  formed  with 
great  evolution  of  heat,  it  is  almost  always  a  very  stable  substance. 

Water  has  the  power  of  combining  with  a  certain  class  of 
chemical  compounds  known  as  the  oxides,  converting  some  of 
them  into  the  important  class  of  compounds  known  as  the  bases, 
and  others  into  the  very  important  class  of  compounds  known 
as  the  acids. 

The  subject  of  acids  and  bases  will  be  considered  in  detail 
when  the  elements  which  form  these  substances  are  studied. 
Many  substances  crystallize  from  solution  in  water  with  a  large 
amount  of  water  in  their  crystals.  This  is  known  as  water  of 
crystallization.  This  water  is  frequently  driven  off  completely 
when  the  crystals  are  heated  to  100°  —  the  boiling-point  of  water. 
In  some  cases  the  crystals  lose  a  part  of  their  water  even  at  ordi- 
nary temperatures.  Such  substances  are  known  as  efflorescent. 

Other  substances  when  exposed  to  the  air  at  ordinary  tempera- 
tures take  up  water.  Such  substances  are  known  as  deliquescent. 

Certain  Physical  Properties  of  Water,  Boiling-point.  —  Water 
at  ordinary  temperatures  and  pressures  is  a  colorless  liquid.  In 
very  thick  layers  it  has  a  bluish  tint.  Under  a  pressure  of 
760  mm.  of  mercury  it  boils  at  100°  C. ;  i.e.  at  this  temperature 
the  tension  of  the  aqueous- vapor  is  just  sufficient  to  overcome 
the  superincumbent  pressure,  which  is  the  pressure  of  the 
atmosphere.  As  the  pressure  to  which  the  water  is  subjected 
increases,  its  boiling-point  rises.  Under  a  pressure  of  five 
atmospheres  the  boiling-point  of  water  is  152°.  Under  a 
pressure  of  ten  atmospheres  water  boils  at  180°.3,  while  under  a 
pressure  of  twenty  atmospheres  it  does  not  boil  until  a  tempera- 
ture of  213°  is  reached. 

These  are  the  conditions  which  obtain  in  a  steam-engine.  The 
water  is  under  the  pressure  of  its  own  vapor,  which  amounts  to 
from  five  to  ten  or  twelve  atmospheres,  and  consequently  its 
boiling-point  is  very  greatly  raised. 


WATER  39 

Similarly,  water  boils  much  lower  on  a  high  mountain  than 
in  a  -valley.  As  we  ascend  a  mountain  the  pressure  of  the 
atmosphere  becomes  less,  and,  consequently,  the  pressure  which 
the  tension  of  the  water-vapor  must  overcome.  On  the  top  of 
Mont  Blanc  water  boils  at  about  84°. 

Heat  of  Vaporization.  —  Any  one  who  ever  observed  water 
boil  must  have  been  impressed  by  the  enormous  amount  of  heat 
which  is  required  to  convert  the  liquid  into  vapor.  He  must 
have  been  further  impressed  by  the  fact  that  the  temperature 
of  the  vapor  is  practically  the  same  as  that  of  the  liquid  from 
which  it  was  formed.  The  amount  of  heat  required  to  convert 
one  gram  of  water  at  100°  into  vapor  at  100°  is  540  calories,  i.e. 
the  same  amount  of  heat  which  would  be  required  to  raise  540 
grams  of  water  1°  in  temperature.  This  is  known  as  the  heat 
of  vaporization  of  water. 

The  Freezing  of  Water.  —  When  water  at  ordinary  pressure  is 
cooled  to  0°  it  freezes,  as  we  say,  or  passes  into  ice.  As  we  cool 
water  down  towards  its  freezing-point,  it  contracts  in  volume 
until  a  temperature  of  4°  is  reached.  As  the  temperature  is 
further  lowered  from  this  point,  the  water  begins  to  expand 
and  continues  to  do  so  until  the  freezing-point  is  reached. 

The  importance  of  this  apparently  insignificant  fact  is  very 
great  indeed  from  the  standpoint  of  the  economy  of  nature. 
Since  water  expands  from  4°  to  the  freezing-point,  ice  is  lighter 
than  water  and  floats  upon  it.  The  importance  of  ice  floating 
upon  water  and  not  sinking  to  the  bottom  is  twofold.  In  the 
first  place,  it  protects  the  water  from  the  extreme  cold  of  the 
atmosphere,  since  ice  is  relatively  a  poor  conductor  of  heat ; 
and  in  the  second  place,  if  ice  sank  to  the  bottom  of  our 
streams  as  fast  as  it  was  formed,  this  would  continually  expose 
a  fresh  surface  of  the  water  to  the  cold,  and  our  streams  might 
be  frozen  solid,  which  would  mean  the  extermination  of  all 
living  things  within  them. 

Heat  of  Fusion  of  Ice.  — We  have  just  learned  what  an  enor- 
mous amount  of  heat  energy  must  be  expended  to  convert  water 
into  vapor.  We  shall  now  see  that  a  large  amount  of  heat  is 


40  ELEMENTS  OF  INORGANIC   CHEMISTRY 

required  to  convert  ice  into  water.  The  amount  of  heat  re- 
quired to  convert  one  gram  of  ice  at  0°  into  water  at  0°  is  80 
calories,  i.e.  the  amount  of  heat  which  would  raise  one  gram 
of  water  from  0°  to  80°.  This  is  known  as  the  heat  of  fusion 
of  ice. 

Of  all  well-known  liquids  water  has  the  highest  specific  heat. 
By  the  specific  heat  of  water  is  meant  the  amount  of  heat 
required  to  raise  a  given  amount  of  water,  say  a  gram,  one 
degree  in  temperature.  This  amount  of  heat,  as  we  have  seen, 
is  one  calorie.  One  calorie  of  heat  will  raise  a  gram  of  any 
other  well-known  substance  more  than  one  degree  in  tempera- 
ture. Water  thus  stands  at  the  head  of  the  list  of  well-known 
substances  as  far  as  specific  heats  are  concerned. 

Similar  results  would  be  obtained  if  we  ran  through  the 
whole  list  of  properties.  Water  would  stand  in  practically 
every  case  either  at  the  top  or  bottom  of  the  list  of  substances. 
Its  properties  are,  therefore,  distinctly  extreme.  They  are 
either  a  maximum  or  a  minimum,  and  usually  a  maximum. 

Solvent  Power  of  Water.  —  Water  has  a  remarkable  power  to 
dissolve  other  substances  which  are  brought  in  contact  with  it. 
Indeed,  of  all  known  substances  it  is  the  best  solvent,  and  with 
respect  to  this  property  it  also  stands  at  the  very  head  of  the 
list  of  chemical  substances.  The  importance  of  solution  for 
chemistry  cannot  be  overestimated.  This  becomes  obvious 
when  we  consider  that  most  chemical  reactions  take  place  in 
solution. 

TInsaturated  and  Saturated  Solutions.  — When  water  has 
dissolved  a  certain  amount  of  a  given  substance,  but  is  still 
capable  of  taking  up  more  of  it,  the  solution  is  said  to  be 
unsaturated.  When  water  has  dissolved  all  of  a  given  sub- 
stance which,  at  the  temperature  in  question,  it  can  take  into 
solution,  the  solution  is  said  to  be  saturated. 


WATER  41 


EXPERIMENTS  WITH  WATER 

Experiment  19.  Distillation;  Separation  of  Volatile  from 
Non-volatile  Substances.  —  (Flask  holding  1  litre  j  condenser ; 
receiver ;  ammonia ;  copper  sulphate.) 

Arrange  an  apparatus  as  in  Figure  15.  Introduce  into  the 
flask  F  water  containing  a  little  ammonia.  Distil  the  liquid. 
Both  water  and  ammonia  will  pass  over,  as  can  be  proved  by 
the  odor  of  the  distillate  and  its  reaction  with  red  litmus. 
This  is  because  ammonia  is  volatile  as  well  as  water. 

Einse  out  the  condenser  and  receiver  and  add  copper  sul- 

Ehate  in  solution  to  the  flask  F,  until  it  no  longer  contains  any 
ree  ammonia.  Now  distil  the  contents  of  the  flask.  Pure 
water  will  pass  over,  since  neither  copper  sulphate  nor  copper 
hydroxide  is  volatile. 

Repeat  the  above  experiment,  adding  to  the  flask  F  water 
containing  a  little  potassium  permanganate.  (See  page  34.) 

Experiment  20.  Electrolysis  of  Water. — (Hofmann  electro- 
lytic apparatus.) 

The  apparatus  used  is  shown  in  Figure  16.  Fill  the  ap- 
paratus with  an  eight  per  cent  solution  of  sulphuric  acid  by 
pouring  the  solution  into  the  large  bulb  R,  the  stop-cocks 
being  open.  As  quickly  as  the  arms  are  entirely  filled,  close 
the  stop-cocks  and  remove  by  means  of  a  piece  of  filter-paper 
any  liquid  in  the  tubes  beyond  the  stop-cocks.  Pass  an 
electric  current  through  the  solution.  The  gases  will  collect 
over  the  two  electrodes  in  the  two  arms  of  the  apparatus.  Over 
one  electrode  the  gas  will  collect  just  twice  as  fast  as  over  the 
other.  After  a  considerable  amount  of  gas  has  been  collected 
in  each  arm,  test  the  smaller  volume  of  gas  for  oxygen  and 
the  larger  volume  for  hydrogen. 

This  experiment  shows  that  two  volumes  of  hydrogen  com- 
bine with  one  volume  of  oxygen,  forming  water.  (See  page  35.) 

Experiment  21.  Quantitative  Synthesis  of  Water  by  the 
Action  of  Hydrogen  on  Copper  Oxide.  —  (Same  apparatus  as 
in  Experiment  13,  with  the  addition  of  a  U-tube.) 

Arrange  an  apparatus  as  in  Figure  17,  conducting  the  hydrogen 
through  a  U-tube  containing  calcium  chloride,  and  a  cylinder 
partly  filled  with  concentrated  sulphuric  acid.  Weigh  the 
amount  of  copper  oxide  used,  and  weigh  the  mixture  of  copper 
and  copper  oxide  after  the  experiment  is  completed.  Weigh 
the  cylinder  containing  concentrated  sulphuric  acid,  before  the 
experiment,  and  weigh  the  cylinder  and  water  after  the  experi- 


42 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


ment.  The  difference  in  weight  is  the  weight  of  the  water 
formed.  Knowing  the  loss  in  the  weight  of  the  copper  oxide, 
we  know  the  weight  of  the  oxygen  required  to  form  the 

T ,     amount  of  water  obtained.     The  difference  between 

the  weight  of  the  water  and  the  weight  of  the  oxygen 
is  the  weight  of  the  hydrogen  required  to  form  the 
water.     It  will  be  found  that  water  is  composed  ap- 
proximately of  one  part  by  weight  of  hydro- 
gen and  eight  parts  by  weight  of  oxygen. 
(See  page  37.) 


FIG.  17. 


Experiment  22.    Efflorescence  of  Sodium  Sulphate.  —  (A  few 

transparent  crystals  of  sodium  sulphate,  and  a  small  beaker  or 
crystallizing  dish.) 

Place  a  few  crystals  of  sodium  sulphate  in  a  small  beaker 
or  crystallizing  dish  and  allow  them  to  stand  exposed  to  the 
air.  'The  crystals  will  lose  water,  become  white  on  the  surface, 
and  opaque.  (See  page  38.) 

Experiment  23.  Deliquescence  of  Calcium  Chloride  and 
Phosphorus  Pentoxide.  —  (Two  small  beakers  or  crystallizing 
dishes;  a  few  lumps  of  calcium  chloride;  a  few  grams  ot 
the  white  powder,  phosphorus  pentoxide.) 

Introduce  the  calcium  chloride  into  one  vessel  and  the  phos- 
phorus pentoxide  into  the  other,  and  expose  both  to  the  air. 
After  a  time  the  calcium  chloride  will  absorb  enough  water 
to  dissolve  it,  and  the  whole  mass  will  become  liquid.  The 
phosphorus  pentoxide  will  absorb  moisture  from  the  air  and 
become  a  gummy  mass.  The  phosphorus  pentoxide  forms 
well-known  compounds  with  water,  as  we  shall  see.  (See  page 
38.) 


CHAPTER   V 

CHLORIN^l 

Occurrence  and  Preparation  of  Chlorine.  —  Chlorine  does  not 
occur  in  the  free  condition  in  nature.  This  is  due  in  part  to 
its  great  chemical  activity.  If  once  set  free  it  would  quickly 
combine  again  with  other  substances.  It  occurs  in  combina- 
tion with  many  other  elements,  but  especially  in  combination 
with  the  element  sodium,  as  sodium  chloride.  The  soluble 
chlorides  do  not  exist  on  the  surface  of  the  earth,  where  they 
are  subject  to  the  influence  of  water,  but  pass  into  solution  and 
are  swept  down  into  the  sea.  This  accounts  for  the  large 
amount  of  chlorine  in  sea-water. 

A  number  of  methods  have  been  devised  for  preparing 
chlorine,  but  most  of  these  are  only  of  historical  interest. 

The  process  devised  by  Deacon  consists  in  oxidizing  hydro- 
chloric acid  by  means  of  the  oxygen  of  the  air.  Hydrochloric 
acid  and  air  are  passed  through  heated  tubes  containing  balls 
of  clay  saturated  with  copper  sulphate.  Under  these  conditions 
the  oxygen  of  the  air  unites  with  the  hydrogen  of  the  hydro- 
chloric acid,  forming  water  and  liberating  chlorine. 

Another  method  for  obtaining  chlorine,  based  upon  the 
oxidation  of  hydrochloric  acid,  is  the  following :  — 

When  a  compound  rich  in  oxygen,  like  manganese  dioxide, 
is  heated  with  hydrochloric  acid,  the  latter  is  oxidized  to  water 
and  chlorine.  A  method  frequently  employed  in  making 
chlorine  in  small  quantities  consists  in  treating  a  mixture 
of  sodium  chloride  and  manganese  dioxide  with  sulphuric 
acid.  Hydrochloric  acid  is  formed  by  the  action  of  the  sul- 
phuric acid  on  the  sodium  chloride,  and  this  is  oxidized  by  the 
manganese  dioxide  liberating  chlorine. 

43 


44  ELEMENTS   OF  INORGANIC   CHEMISTRY 

A  very  convenient  method  for  obtaining  chlorine  in  the  lab- 
oratory consists  in  treating  bleaching-powder  with  hydrochloric 
acid.  The  bleaching-powder  is  introduced  into  an  ordinary  Kipp 
apparatus,  and  the  acid  allowed  to  come  in  contact  with  it. 

All  of  these  methods  have  practically  given  place  to  the 
electrolytic,  when  chlorine  is  to  be  prepared  on  the  large  scale. 
Most  of  the  chlorine  is  now  prepared  by  the  electrolysis  of 
aqueous  potassium  or  sodium  chloride.  When  a  solution  of 
potassium  chloride  in  water  is  electrolyzed,  hydrogen  separates 
at  one  pole,  and  chlorine  at  the  other.  The  potassium  remains 
in  solution  as  potassium  hydroxide. 

Chemical  Properties  of  Chlorine.  —  The  yellowish  green  gas 
chlorine  is  chemically  very  active.  It  combines  with  many 
elements  by  simple  contact,  and  often  with  evolution  of  much 
heat  and  even  light,  and  in  some  cases  with  explosive  violence. 

When  copper  foil  is  brought  in  contact  with  chlorine  gas,  it 
combines  with  the  chlorine,  shown  by  the  fact  that  it  glows 
and  forms  chloride  of  copper. 

When  finely  divided  antimony  is  allowed  to  fall  into  a  vessel 
containing  chlorine  gas,  we  have  literally  a  rain  of  fire  —  each 
antimony  particle  becoming  incandescent  as  it  combines  with 
the  chlorine.  Other  substances,  like  brass,  burn  in  chlorine 
only  when  they  have  been  heated  to  an  elevated  temperature. 

Combustion  in  Chlorine.  —  We  have  here  examples  of  combi- 
nation taking  place  between  substances  and  chlorine,  which  are 
analogous  to  combustion  in  oxygen.  In  the  former  case,  as  in 
the  latter,  the  combination  takes  place  with  evolution  of  light 
and  heat,  and  the  combustion  in  chlorine  is  even  more  energetic 
than  in  oxygen,  in  that  it  starts  at  ordinary  temperatures.  We 
have,  then,  combustion  in  chlorine  just  as  truly  as  in  oxygen. 
The  term  "  combustion,  "  however,  as  ordinarily  used  always  re- 
fers to  combination  with  oxygen,  since  we  never  know  chlorine 
in  the  free  condition  unless  it  is  specially  prepared. 

Action  of  Chlorine  on  Hydrogen. — Hydrogen  unites  with 
chlorine  at  ordinary  temperatures  if  exposed  to  diffuse  light, 
and  with  explosive  violence  if  exposed  to  direct  sunlight.  A 


CHLORINE  45 

jet  of  hydrogen  can,  however,  be  burned  in  chlorine  just  as  it 
can  be  burned  in  oxygen,  and  a  jet  of  chlorine  can  be  burned 
in  hydrogen.  When  hydrogen  is  burned  in  oxygen,  the  two 
gases  combine,  forming  water.  When  hydrogen  is  burned  in 
chlorine  the  two  gases  combine,  forming  the  important  com- 
pound hydrochloric  acid,  which  we  shall  study  a  little  later. 

Action  of  Chlorine  on  Water.  —  Chlorine  is  readily  soluble 
in  water,  and  the  resulting  solution  is  known  as  chlorine  ivater. 
Chlorine  water,  if  kept  in  the  dark,  is  a  stable  substance,  but  if 
exposed  to  the  light,  a  deep-seated  change  takes  place.  The 
chlorine  acts  chemically  upon  the  water,  combining  with  the 
hydrogen  and  liberating  oxygen.  The  resulting  solution  con- 
tains the  hydrochloric  acid  formed,  while  the  oxygen  gas  is 
liberated.  Such  chemical  reactions  which  are  brought  about 
by  the  action  of  light  are  known  as  photochemical  reactions. 
Since  oxygen  is  liberated,  chlorine  is  known  as  a  strong  oxidiz- 
ing agent.  Its  oxidizing  power  renders  chlorine  one  of  the  very 
best  bleaching  agents  which  is  at  our  disposal.  The  oxygen 
which  is  set  free  when  chlorine  acts  on  moisture  oxidizes 
organic  coloring-matter,  and  leaves  behind  a  colorless  sub- 
stance. This  can  be  illustrated  by  bringing  into  the  presence 
of  chlorine  gas  some  moist  flowers  or  a  moist  piece  of  calico, 
when  the  color  will  disappear  in  a  very  short  time. 

Chlorine  is  also  an  excellent  disinfectant,  having  an  unusual 
power  to  destroy  bacteria  and  other  forms  of  life.  This  is  due 
in  part  to  its  oxidizing  action,  and  in  part  to  direct  combination 
of  chlorine  with  the  organic  matter  of  such  forms  of  life. 

Action  of  Chlorine  on  Certain  Organic  Compounds. — Chlorine 
not  only  acts  on  elementary  substances  and  simple  compounds, 
but  also  on  complex  organic  substances.  When  brought  in 
contact  with  organic  substances  which  contain  hydrogen,  the 
chlorine  first  replaces  the  hydrogen,  taking  its  place  in  the 
molecule,  and  then  more  chlorine  combines  with  the  replaced 
hydrogen,  forming  hydrochloric  acid.  Such  a  reaction,  known 
as  substitution,  can  be  illustrated  by  bringing  a  piece  of  filter- 
paper  saturated  with  oil  of  turpentine  into  a  vessel  filled  with 


46  ELEMENTS   OF   INORGANIC   CHEMISTRY 

chlorine  gas.  A  violent  reaction  takes  place,  resulting  in  the 
liberation  of  a  large  amount  of  finely  divided  carbon. 

Chlorine  Hydrate.  —  When  chlorine  gas  is  conducted  into  a 
mixture  of  water  and  ice,  a  crystalline  compound  separates. 
At  ordinary  temperatures  it  decomposes  into  chlorine  and 
water,  while  at  somewhat  elevated  temperatures  the  decompo- 
sition is  quite  rapid,  resulting  in  a  copious  evolution  of  chlorine 
gas.  This  compound  is  of  special  historical  interest  in  connec- 
tion with  the  liquefaction  of  chlorine.  The  earlier  work  on 
the  liquefaction  of  gases  was  carried  out  almost  exclusively  by 
the  great  English  physicist,  Faraday.  He  succeeded  in  liquefy- 
ing chlorine  obtained 
from  chlorine  hydrate. 
Some  of  this  com- 
pound was  placed  in 
one  end  of  a  thick- 
walled  glass  tube,  and 
the  other  end  closed, 
as  shown  in  Figure 
18.  The  end  of  the 

tube  containing  the  chlorine  hydrate  was  gently  warmed,  while 
the  other  end  was  surrounded  by  a  freezing-mixture  of  ice  and 
salt.  Under  these  conditions  the  chlorine  was  liberated,  and 
produced  a  pressure  in  the  tube  which  was  sufficient  to  liquefy 
it.  Considerable  precaution  is  necessary  in  carrying  out  this 
experiment,  since  an  explosion  is  liable  to  result. 

Certain  Physical  Properties  of  Chlorine.  —  The  yellowish 
green  gas,  chlorine,  is  about  two  and  one-half  times  as  heavy 
as  the  air,  a  litre  weighing  3.22  grams.  It  has  a  most  disagree- 
able odor,  and  an  injurious  effect  when  inhaled.  It  acts  upon 
the  mucous  membrane  of  the  nose  and  throat,  and  disintegrates 
these  tissues  if  inhaled  in  sufficient  quantity  and  for  sufficient 
time.  It  is  therefore  necessary  in  working  with  chlorine  to 
take  every  precaution  to  be  protected  from  the  gas.  Such 
work  should  always  be  done  under  a  good  hood,  with  a  strong 
draft  to  remove  the  gas  as  rapidly  as  it  escapes  into  the  air. 


CHLORINE 


47 


Comparative  Inactivity  of  Dry  Chlorine.  —  While  moist 
chlorine  is  one  of  the  most  active  substances  chemically,  dry 
chlorine  is  comparatively  inactive.  A  lecture  table  experi- 
ment, which  is  frequently  shown,  is  to  pass  chlorine  through  a 
glass  tube  containing  a  piece  of  metallic  sodium  which  is  heated 
by  means  of  a  bunsen  burner.  If  the  chlorine  has  been  care- 
fully dried,  as  is  frequently  done,  the  sodium  will  melt  and  re- 
main with  untarnished  surface  in  contact  with  the  chlorine  gas. 
If,  on  the  other  hand,  the  water-vapor  has  not  been  removed 
from  the  chlorine,  vigorous  chemical  action  will  take  place, 
accompanied  by  intense  heat  and  a  bright  light,  and  the  com- 
pound sodium  chloride  will  be  formed. 


EXPERIMENTS  WITH  CHLORINE 


Experiment  24.  Prepa- 
ration of  Chlorine  from 
Hydrochloric  Acid  and  Man- 
ganese Dioxide.  —  (A  two- 
litre  balloon  flask;  safety- 
tube  ;  rubber  stopper ; 
manganese  dioxide ;  hydro- 
chloric acid.) 

Arrange  an  apparatus  as 
in  Figure  19.  Into  the  flask 
A  introduce  some  coarse- 
grained manganese  diox- 
ide. Pour  hydrochloric  acid 
through  the  safety-tube  un- 
til the  dioxide  is  covered 
with  the  liquid.  -  Heat 
gently  ;  chlorine  is  set  free. 
The  gas  may  be  passed 
through  water  in  a  cylinder 
to  remove  all  traces  of  hy- 
drochloric acid,  and  can  be 
collected  from  the  tube. 
This,  and  all  experiments 
with  chlorine,  should  be 
performed  under  a  good 


FIG.  19. 


48 


ELEMENTS   OF   INORGANIC   CHEMISTRY 


hood  with  a  strong  draft.     Any  chlorine  not  collected  should 
be  conducted  into  a  solution  of  caustic  potash.     (See  page  43.) 
Experiment  25.     Preparation  of  Chlorine  by  the  Action  of 
Hydrochloric  Acid  on  Potassium  Bichromate.  —  (Same  apparatus 
as  in  last  experiment.    In  addition  some  potassium  dichromate.) 
Arrange  the  apparatus  exactly  as  in  the  preceding  experi- 
ment.    Introduce  into  the  flask  A  50  or  60  grams  of  potas- 
sium  dichromate  which  has  been  coarsely 
pulverized.    Pour  into  the  flask  through  the 
thistle  tube  200  to  250  grams  of  concentrated 
hydrochloric  acid.     Warm  gently  the  con- 
tents    of    the     flask.     An    abundance     of 
pure    chlorine    can    be    obtained    by   this 
method. 

Experiment  26.    Prepara- 
tion of  Chlorine  from  Sodium 
Chloride,  Manganese  Dioxide, 
and  Sulphuric  Acid.  —  (Same 
apparatus    as    in    the    pre- 
ceding experiment;  sodium 
chloride ;  sulphuric  acid  con- 
taining two   parts    by 
weight  of  acid  to  one 
of  water.) 

Arrange  an  appara- 
tus as  in  Figure  20. 
Introduce  into  the  flask 
A  a  mixture  of  50 
grams  of  sodium  chlo- 
ride and  50  grams  of 
finely  powdered  man- 
ganese dioxide.  Pour 
into  the  flask  through 
the  safety-tube  50  cc.  of  the  sulphuric  acid,  obtained  by  adding 
two  parts  by  weight  of  the  concentrated  acid  to  one  of  water 
and  allowing  the  mixture  to  cool.  When  gently  heated  an 
abundance  of  chlorine  is  obtained. 

Fill  a  number  of  cylinders  or  wide-mouthed  bottles  with  the 
gas,  by  allowing  the  tube  from  which  the  chlorine  escapes  to 
reach  to  the  bottom  of  the  receiving  vessel.  The  chlorine 
being  heavier  than  air  will  displace  the  latter  and  fill  the 
vessel.  It  is  well  to  cover  the  vessel  which  is  to  be  filled  with 
chlorine  with  a  perforated  glass  plate  through  which  the  de- 


FIG.  20. 


CHLORINE 


49 


livery  tube  is  passed.  In  this  way  air  currents  are  avoided, 
and  the  vessel  can  be  more  perfectly  filled  with  the  pure 
chlorine.  The  vessels  thus  filled  with  the  gas  should  be  set 
aside  for  use  in  the  following  experiments.  (See  page  43.) 

Experiment  27.  Bleaching  Action  of  Chlorine.  —  (Vessel 
filled  with  chlorine  ;  flowers ;  colored  fabric ;  etc.) 

Moisten  the  flowers,  or  the  fabric,  and  introduce  into  the 
chlorine.  After  a  time  the  color  will  be  practically  destroyed. 
(See  page  45.) 

Experiment  28.  Combination  of  Chlorine  with  the  Metals.  — 
(Three  vessels  filled  with  chlorine ;  thin  copper  foil ;  Dutch 
leaf ;  powdered  antimony.) 

Into  one  vessel  filled  with  chlorine  introduce  the  copper  foil, 
into  another  the  antimony  powder,  and  into  the  third  Dutch 
leaf.  There  is  abundant  evidence  of  the  combination  of  the 
chlorine  with  the  metals.  (See  page  44.) 

Experiment  29.  Action  of  Chlorine  on  Organic  Compounds 
containing  Hydrogen.  —  (Vessel  filled  with  chlorine ;  filter-paper ; 
oil  of  turpentine.) 

Dip  a  strip  of  filter-paper  into  warm  oil  of  turpentine  and 
introduce  into  a  vessel  filled  with  chlorine.  A  vigorous  reac- 
tion will  take  place,  resulting  in  the  liberation  of  a  large  amount 
of  carbon  which  will  cover  the  walls 
of  the  vessel.  (See  page  45.) 

Experiment  30.  Action  of  Light 
on  Chlorine  Water.  —  (Long  glass 
tube  2  to  2i  cm.  in  diameter ;  chlo- 
rine water.) 

Fill  a  glass  tube  with  chlorine 
water,  obtained  by  saturating  cold 
water  with  chlorine.  Expose  the 
tube  to  diffuse  light  and  to  ordinary 
sunlight  and  observe  the  bubbles  of 
oxygen  rise  in  the  tube  and  collect 
at  the  top.  (See  page  45.) 

Experiment  31.  Burning  Hydro- 
gen in  Chlorine.  —  (Cylinder  filled 
with  chlorine ;  hydrogen  generator.) 

Arrange  the  apparatus  as  shown  in  Figure  21.  The  hydrogen 
escaping  from  the  tube  is  ignited  and  lowered  into  the  chlorine. 
The  hydrogen  burns  with  a  light  blue  flame,  and  vapors  of 
hydrochloric  acid  escape.  This  can  be  seen  at  the  mouth 
of  the  cylinder,  or  tested  by  a  piece  of  moistened  blue  litmus 
paper,  which  is  turned  red.  (See  page  44.) 


FIG.  21. 


CHAPTER  VI 

LAWS     OF     CHEMICAL    ACTION 

THE  study  of  chemical  phenomena  consisted  at  first  in 
simple  observation  and  description.  Two  substances  were 
allowed  to  react  chemically  and  the  reaction  was  observed. 
The  nature  and  properties  of  the  substances  entering  into 
the  reaction  were  studied,  and  then  the  nature  and  properties 
of  the  products  formed.  This  was  the  qualitative  stage  of 
chemistry. 

Mere  qualitative  observations  are  followed  by  quantitative 
measurements  in  the  development  of  any  branch  of  natural 
science. 

The  quantities  of  the  substances  which  enter  into  a  reaction 
were  carefully  weighed,  and  the  quantities  of  the  products 
formed.  The  amounts  of  different  substances  which  combine 
with  one  another  were  determined,  and  certain  other  changes 
which  take  place  simultaneously  with  chemical  transformations 
were  studied.  Still  we  do  not  have  a  science.  Indeed,  the 
highest  aim  of  a  science  is  not  simply  to  observe  and  record 
facts. 

Generalization.  —  The  highest  aim  of  scientific  investigation 
is  the  discovery  of  wide-reaching  relations  between  large  num- 
bers of  facts.  Such  relations  when  sufficiently  comprehensive 
are  known  as  generalizations.  A  simple  example  by  way  of 
illustration.  It  was  early  observed  that  a  body  thrown  upward 
from  the  surface  of  the  earth  will  return  again  to  the  surface. 
Repeated  observations  confirmed  those  first  made,  but  it 
remained  for  Newton  to  arrive  at  the  generalization  known  as 
the  law  of  gravitation. 

50 


LAWS  OF   CHEMICAL  ACTION  51 

In  a  similar  manner  certain  generalizations  have  been 
reached  in  chemistry,  which  have  been  of  fundamental  impor- 
tance in  the  development  of  the  science.  Some  of  these  will 
be  considered  in  this  place,  while  others  will  be  introduced  in 
connections  into  which  they  seem  to  enter  naturally. 

The  Law  of  the  Conservation  of  Mass.  —  We  have  seen  that 
when  substances  react  chemically  they  disappear  as  such,  and 
products  are  formed  having  properties  very  different  from  the 
original  substances.  When  copper  is  heated  in  the  air,  a  black 
powder  is  formed  having  properties  which  are  very  different 
from  the  original  copper.  The  question  arises,  Are  all  the 
properties  of  the  copper  lost  during  the  chemical  transforma- 
tion, or  have  only  some  of  them  disappeared  ?  Take  the  prop- 
erty mass.  Does  the  mass  of  the  substances  entering  into  a 
chemical  reaction  undergo  any  change  during  the  reaction  ? 

Since  weight  is  a  measure  of  mass,  the  problem  reduces 
itself  to  determining  whether  there  is  any  change  in  weight 
when  chemical  reaction  takes  place. 

It  has  been  found  that  there  is  no  appreciable  change  in 
weight  in  chemical  reaction,  consequently,  no  appreciable 
change  in  mass,  and  this  is  known  as  the  law  of  the  conserva- 
tion of  mass. 

The  Law  of  Constant  Proportion.  —  The  second  important 
generalization  which  was  reached  through  the  quantitative 
study  of  chemical  phenomena,  was  that  the  constituents  of  a 
chemical  compound  are  always  present  in  a  constant  propor- 
tion. If  two  substances  react  chemically  and  form  a  third, 
they  enter  into  combination  in  a  constant  proportion.  The  law 
may  be  formulated  thus  :  — 

Any  given  chemical  compound  always  contains  the  same  con- 
stituents, and  there  is  a  constant  proportion  between  the  masses  of 
the  constituents  present. 

The  Law  of  Multiple  Proportions.  —  While  it  is  true  that 
substances  combine  in  constant  proportions,  it  is  also  true  that 
two  substances  may  combine  in  more  than  one  proportion. 


52  ELEMENTS   OF  INORGANIC   CHEMISTRY 

The  two  compounds,  methane  (CH4)  and  ethylene  (C2H4), 
were  analyzed  and  it  was  found  that  the  ratio  of  carbon  to 
hydrogen  in  the  former  was  as  3  to  1 ;  in  the  latter  as  6  to  1. 
The  latter  evidently  contains  twice  as  much  carbon  with  respect 
to  hydrogen  as  the  former.  A  number  of  analogous  cases 
where  two  elements  combine  in  more  than  one  proportion 
were  examined,  and  the  result  was  the  discovery  by  Dalton  of 
the  law  of  multiple  proportions.  This  law  may  be  formulated 
thus : — 

If  two  elements  combine  in  more  than  one  proportion,  the 
masses  of  the  one  which  combine  with  a  given  mass  of  the  other 
bear  a  simple  rational  relation  to  one  another. 

Since  this  law  was  proposed,  great  masses  of  facts  which 
bear  upon  it  have  been  discovered.  The  result  is  that  the  law 
has  been  found  to  hold  thus  far  without  an  exception.  Dalton 
raised  the  question,  What  do  the  laws  of  definite  and  multiple 
proportions  really  mean?  Why  do  such  relations  obtain? 
His  answer  is  what  has  come  to  be  known  as  the  atonyc  theory. 
The  view  that  matter  is  composed  of  indivisible  particles  or 
atoms,  which  have  definite  weights,  and  that  chemical  action 
takes  place  between  these  particles,  was  to  Dalton  the  only 
rational  explanation  of  the  laws  of  multiple  proportion  and 
combining  weights.  If  water  is  composed  of  such  ultimate, 
indivisible  parts  or  atoms,  then  a  constant  number  of  atoms  of 
one  substance  combines  with  one  atom  of  another  substance  to 
form  a  definite  molecule  of  the  compound,  and  we  have  the  law 
of  constant  proportions.  One  atom  of  one  substance  may 
combine  with  one  atom  of  another  substance,  or  a  number  of 
atoms  of  one  substance  may  combine  with  one  of  another  to 
form  a  molecule,  but  the  number  must  be  a  simple,  rational, 
whole  number  ;  whence  the  law  of  multiple  proportions. 

The  Correlation  and  Conservation  of  Energy.  —  We  have  been 
considering  thus  far  entirely  the  material  transformations 
that  take  place  in  chemical  reactions,  and  have  pointed  out 
certain  generalizations  which  have  been  reached,  and  which 


LAWS   OF   CHEMICAL  ACTION  53 

lie  at  the  foundation  of  the  science  of  chemistry.  Were  we  to 
stop  here  we  should  leave  untouched  a  class  of  phenomena 
whose  importance  cannot  easily  be  overestimated. 

Whenever  we  have  chemical  reaction  taking  place  we  have 
heat  liberated  or  absorbed,  and  usually  liberated.  This  brings 
us  to  a  study  of  the  energy  changes  which  are  inseparably  con- 
nected with  all  chemical  action. 

Energy  manifests  itself  in  a  number  of  forms.  We  have 
light  energy,  heat  energy,  electrical  energy,  mechanical  energy, 
and  these  are  mutually  convertible  into  one  another.  That 
mechanical  energy  can  be  converted  into  heat  is  shown  wher- 
ever friction  exists.  Rub  together  two  pieces  of  metal  and 
both  become  hot.  That  heat  energy  can  be  converted  into 
light  energy  is  illustrated  by  a  piece  of  metal  which  has  been 
heated  to  incandescence.  This  principle  of  the  mutual  con- 
vertibility of  the  various  forms  of  energy  is  known  as  the 
principle  of  the  correlation  of  energy,  and  is  an  important 
generalization  in  physical  science. 

That  one  form  of  energy  can  be  converted  qualitatively  into 
another  is  important,  but  far  less  important  than  the  fact  that 
one  form  of  energy  can  be  converted  quantitatively  into  another. 
When,  for  example,  mechanical  energy  disappears,  as  when  a 
hammer  falls  upon  a  metal  plate,  the  heat  energy  produced  is 
exactly  equivalent  to  the  mechanical  energy  which  has  dis- 
appeared. If  the  heat  energy  produced  under  these  conditions 
was  transformed  into  work,  it  would  raise  the  hammer  again 
exactly  to  its  original  position.  This  principle,  fundamental 
to  the  science  of  physics,  is  known  as  the  principle  of  the  con- 
servation of  energy.  It  says  in  words  that  no  energy  can  be 
created  or  lost,  and  is  analogous  to  the  law  of  the  conservation 

»of  mass,  which  we  have  already  studied. 
Importance  of  the  Conservation  of  Energy  for  the  Science  of 
Chemistry.  —  The  bearing  of  the  conservation  of  energy  upon 
chemistry  may  not  appear  at  first  sight.  In  addition  to  the 
forms  of  energy  enumerated  above  we  should  add  intrinsic 
energy,  which  is  frequently  referred  to  as  chemical  energy.  This 


54  ELEMENTS  OF   INORGANIC   CHEMISTRY 

form  of  energy  exists  in  all  substances  in  larger  or  smaller 
amounts,  and  is  the  form  which  is  converted  into  heat  when  a 
piece  of  coal  is  burned.  The  existence  of  this  form  of  energy 
is  essential  to  all  chemical  action,  and  is,  therefore,  absolutely 
essential  to  the  science  of  chemistry. 

Although  only  a  part  of  the  intrinsic  or  chemical  energy  in 
the  substances  which  react  is  converted  into  heat  or  some  other 
form  of  energy  during  the  reaction,  yet  this  part  which  dis- 
appears is  converted  quantitatively  into  other  forms.  The  law 
of  the  conservation  of  energy  is,  therefore,  fundamental  to  the 
scientific  study  of  chemistry. 

Chemical  Formulas  and  Chemical  Equations.  —  The  formula 
for  any  chemical  compound  is  the  symbols  of  the  elements  in 
the  compound  written  side  by  side.  Thus,  the  formula  for 
water  is  H20.  This  means  that  the  molecule,  or  the  smallest 
particle  of  the  compound  of  water,  is  made  up  of  two  atoms  of 
hydrogen  and  one  atom  of  oxygen.  The  formula  for  manga- 
nese dioxide  is  Mn02,  which  means  that  the  molecule  or  small- 
est particle  of  the  compound  contains  one  atom  of  manganese 
and  two  atoms  of  oxygen.  The  formula  for  potassium  chlorate 
is  KC103,  the  molecule  containing  one  atom  of  potassium,  one 
atom  of  chlorine,  and  three  atoms  of  oxygen. 

Chemical  reactions  are  expressed  by  chemical  equations,  as 
we  say.  The  term  "  equation  "  is  used  because  the  same  kinds 
and  amounts  of  matter  must  appear  on  both  sides  of  it. 

Take  the  reaction  between  hydrogen  and  chlorine.  This  is 
expressed  by  the  following  equation :  — 

1H2  +  C12  =  2HC1, 

which  means  that  two  atoms  of  hydrogen  unite  with  two  atoms 
of  chlorine  and  form  two  molecules  of  hydrochloric  acid. 

1  It  is  impossible  to  explain  until  later  the  reason  for  using  the  apparent 
double  quantities  of  the  elements  in  this  and  subsequent  equations.  It  is, 
however,  better  to  learn  the  true  equations  at  once,  even  if  their  full 
significance  cannot  be  seen  until  a  little  more  progress  has  been  made  in 
the  subject. 


LAWS  OF   CHEMICAL  ACTION  55 

The  reaction  between  hydrogen  and  oxygen  is  expressed  by 
the  following  equation  :  — 

2  H2  +  02  =  2  H20. 
i 

This  means  that  four  atoms  of  hydrogen  unite  with  two  atoms 
of  oxygen  and  form  two  molecules  of  water. 

In  a  similar  manner  chemical  reactions  in  general  can  be 
expressed  by  equations,  which,  in  some  cases,  are  quite  com- 
plex. We  can  now  write  the  equations  for  the  reactions 
already  studied.  The  equation  in  any  case  does  not  tell  us 
simply  what  elements  enter  into  a  reaction,  but  the  quantities 
of  the  different  elements  that  react,  and  the  number  of  atoms 
of  the  different  elements  in  the  various  substances  involved  in 
the  reaction.  This  will  be  seen  in  the  following  examples. 

Equations  involved  in  the  Study  of  Oxygen.  —  The  formation 
of  oxygen  from  potassium  chlorate  may  be  represented  by  the 
following  equation  :  — 

2KC103=2KC1  +  302. 

When  oxygen  is  formed  by  heating  manganese  dioxide,  we 
have  :  — 

3  Mn02  =  Mn304  +  02. 

The  formation  of  oxygen  from  mercuric  oxide  is  represented 
thus  :  — 

2HgO  =  2Hg  +  02; 

while  the  reaction  between  hydrogen  dioxide  and  lead  dioxide 
may  be  formulated  thus  :  — 

H202  +  Pb02  =  PbO  +  H20  +  02. 
When  sulphur  burns  in  oxygen,  we  have  :  — 


The  combustion  of  iron  is  represented  thus  :  — 


56  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Carbon  burns  in  oxygen,  forming  carbon  dioxide  :  — 

C  +  02=C02; 

while  phosphorus  forms  the  pentoxide  :  — 
4P  +  502  =  2P205. 

Equations  involved  in  the  Study  of  Hydrogen.  —  The  follow- 
ing equation  represents  the  reaction  between  sodium  and 
water  :  — 

2  Na  +  2  H20  =  2  NaOH  +  H2. 

When  hydrochloric  acid  acts  on  zinc,  we  have  :  — 

2  HC1  +  Zn  =  ZnCl2  +  H2. 
When  sulphuric  acid  acts  on  zinc  :  — 

H2S04  +  Zn  =  ZnS04  +  H2. 

The  reaction  between  hydrogen  and  oxygen  may  be  rep- 
resented thus  :  — 

=  2H20. 


The  reduction  of  copper  oxide  by  hydrogen  is  formulated  as 
follows  :  — 

CuO-f  H2  = 


The  decomposition  of  water  by  the  electric  current  can  be 
represented  thus  :  — 


Equations  involved  in  the  Study  of  Chlorine.  —  When  hydro- 
chloric acid  is  oxidized  by  the  oxygen  of  the  air  :  — 


When  it  is  oxidized  by  manganese  dioxide  :  — 

4  HC1  +  Mn02  =  2  H20  +  MnCl2  +  C12. 


LAWS   OF  CHEMICAL.  ACTION  57 

The  reaction-  when  sodium  chloride  is  treated  with  sulphuric 
acid  in  the  presence  of  manganese  dioxide  is  more  complicated 
and  may  be  represented  as  taking  place  in  two  stages  :  — 

I.     2  NaCl  4-  H2S04  =  Na2S04  +  2  HC1. 
II.     2  HC1  +  Mn02  +  H2S04  =  MnS04  +  2  H20  +  C12. 

Eepresenting  both  reactions  in  one  equation,  we  have  :  — 
2  NaCl  +  Mn02  +  2  H2S04  =  MnS04  +  Na2S04  +  2  H20  +  C12. 


The  reaction  between  hydrochloric  acid  and  potassium  di- 
chromate,  giving  free  chlorine,  may  be  formulated  thus  :  — 

K2Cr207  +  14  HC1  =  2  KC1  +  2  CrCl,  +  7  H20  +  3  C12. 

The  reaction  between  copper  and  chlorine  may  be  represented 
thus  :  — 

Cu  +  C12  =  CuCl2. 

Between  antimony  and  chlorine  thus  :  — 
2  Sb  +  3  C12  =  2  SbCl3. 

The  combination  of  chlorine  and  hydrogen  is  represented 
very  simply  :  — 

H2  +  C12  =  2  HC1  ; 

while  the  reaction  of  chlorine  with  water  is  a  little  more 
complex  :  — 

2  C12  +  2  H20  =  4  HC1  +  02. 

The  action  of  chlorine  on  organic  compounds  is  much  more 
complicated. 


CHAPTER   VII 

OZONE   AND   HYDROGEN   DIOXIDE 
OZONE 

Allotropic  Modification  of  Oxygen.  —  We  have  dealt  thus  far 
with  the  element  oxygen  in  the  condition  in  which  it  is  ordi- 
narily known  to  us.  Oxygen  can,  however,  occur  with  prop- 
erties very  different  from  ordinary  oxygen.  The  second 
modification  of  oxygen  is  known  as  ozone.  "The  property  of 
an  element  to  occur  in  two  different  modifications  is  known  as 
allotropy,  and  ozone  is  spoken  of  as  an  allotropic  modification 
of  oxygen?^) 

Preparation  of  Ozone.  —  Ozone  is  formed  in  larger  or  smaller 
quantities  under  a  number  of  conditions.  When  phosphorus 


FIG.  22.    ' 

is  exposed  to  the  air  it  undergoes  slow  oxidation,  and  at  the 
same  time  some  of  the  oxygen  of  the  air  is  converted  into 
ozone. 

The  best  method  of  obtaining  ozone  in  quantity  is  by  pass- 
ing electricity  through  oxygen. 

A  convenient  form  of  apparatus  for  preparing  ozone  is  the 
following  (Fig.  22) :  Into  the  glass  tube  GG  an  iron  tube  77  is 
inserted.  The  glass  tube  is  surrounded  for  a  part  of  its  length 

58 


OZONE  AND   HYDROGEN  DIOXIDE  59 

by  tin-foil.  Oxygen  is  introduced  into  the  glass  tube  through 
the  tube  A  and  escapes  through  B.  A  current  of  water  is 
passed  through  the  tube  (7(7  to  keep  the  apparatus  cool.  The 
tin-foil,  on  the  one  hand,  and  the  tube  (7,  on  the  other,  are 
connected  with  the  poles  of  an  induction  machine.  Silent 
discharges  take  place  between  the  tin-foil  and  the  iron,  passing 
through  the  oxygen.  Under  these  conditions  a  part  of  the 
oxygen  is  converted  into  ozone. 

Transformation  of  Ozone  into  Oxygen.  —  We  have  seen  that 
oxygen  is  transformed  into  ozone  under  the  influence  of  the 
silent  electrical  discharge.  The  question  naturally  arises,  Can 
ozone  once  formed  be  transformed  again  into  oxygen?  The 
answer  is,  it  can.  When  ozone  is  heated  to  300°,  it  passes  back 
into  ordinary  oxygen.  We  can  thus  pass  either  from  oxygen 
to  ozone  or  from  ozone  to  oxygen. 

This  raises  the  important  question,  What  is  the  cause  of  the 
difference  in  properties  between  the  two  modifications  of 
oxygen  ? 

The  Same  Kind  of  Matter  but  Different  Amounts  of  Energy. 
—  We  have  already  seen  that  oxygen  and  ozone  are  made  up 
of  the  same  kind  of  matter,  since  each  is  transformable  into 
the  other.  If  we  study  this  side  of  the  problem  quantitatively, 
we  shall  find  that  when  three  volumes  of  oxygen  are  converted 
into  ozone,  the  resulting  gas  occupies  only  two  volumes.  Thus, 
if  three  litres  of  oxygen  were  converted  into  ozone,  only  two 
litres  of  ozone  would  be  formed.  On  the  other  hand,  if  two 
litres  of  ozone  were  decomposed  by  heat,  three  litres  of  oxygen 
would  be  formed. 

To  anticipate  what  we  shall  understand  more  clearly  later, 
the  atom  of  oxygen  cannot  exist  by  itself  in  the  free  state,  but 
two  atoms  of  oxygen  always  unite  and  form  what  is  called  the 
molecule  of  oxygen.  In  oxygen  gas,  as  we  ordinarily  know  it, 
we  do  not  have  the  atoms  of  oxygen  uncombined  with  one 
another,  but  the  molecules  which  are  formed  by  the  union  of 
two  atoms. 

It  has  been  shQwji  that  in  the  niglecule  of  ozone  there  are 


60  ELEMENTS  OF  INORGANIC   CHEMISTRY 

three  atoms  of  oxygen,  while  in  the  molecule  of  oxygen  there 
are  only  two.  It  is  obvious,  however,  that  the  difference  in 
the  number  of  atoms  in  the  molecule,  alone  considered,  is  not 
sufficient  to  account  for  such  differences  in  properties  as  exist 
between  oxygen  and  ozone.  Indeed,  it  is  difficult  to  see  how 
this  would  produce  a  difference  in  any  property  other  than  the 
mass  of  the  molecule. 

The  real  difference  in  the  properties  of  oxygen  and  ozone  is 
due  to  the  different  amounts  of  intrinsic  energy  present  in 
their  molecules.  This  statement  is  not  made  dogmatically, 
but  can  be  demonstrated  experimentally  in  the  following 
manner :  — 

When  carbon  is  burned  in  oxygen  the  product  is  carbon 
dioxide.  When  carbon  is  burned  in  ozone  the  product  is 
carbon  dioxide.  We  start  with  the  same  substance,  carbon, 
in  both  cases,  and  we  end  with  the  same  product,  carbon 
dioxide.  Any  differences  in  the  two  reactions  must  be  due  to 
the  differences  between  the  oxygen  and  the  ozone. 

If  we  measure  the  amounts  of  heat  liberated  in  the  two 
reactions,  we  find  that  they  are  very  different  indeed.  Con- 
siderably more  heat  is  evolved  when  carbon  is  burned  in  ozone 
than  when  carbon  is  burned  in  oxygen.  This  shows  that 
there  is  more  intrinsic  energy  present  in  the  molecule  of  ozone 
than  in  the  molecule  of  oxygen. 


HYDROGEN  DIOXIDE 

Hydrogen  Dioxide.  —  One  compound  of  hydrogen  and  oxy- 
gen, other  than  water,  calls  for  special  comment.  This  is  the 
compound  hydrogen  dioxide.  It  has  the  composition  expressed 
by  the  formula  H202,  and  is,  therefore,  an  oxidized  water. 
Hydrogen  dioxide  is  most  readily  prepared  in  quantity  by 
treating  the  compound  barium  dioxide,  with  an  acid.  If  we 
use  hydrochloric  acid,  the  equation  is  expressed  thus :  — 

Ba02  +  2  HC1  =  BaCl2  +  H202. 


OZONE   AND   HYDROGEN  DIOXIDE  61 

Properties  of  Hydrogen  Dioxide.  —  Hydrogen  dioxide  is  a 
colorless,  viscous  liquid,  much  heavier  than  water,  having  a 
specific  gravity  of  1.4996.  We  saw  that  water  in  thick  layers 
has  a  markedly  bluish  tint.  Hydrogen  dioxide  is  still  deeper 
blue  when  observed  in  thick  layers. 

One  of  the  characteristic  properties  of  hydrogen  dioxide  is 
the  ease  with  which  it  decomposes  into  water  and  oxygen. 


On  account  of  the  ease  with  which  hydrogen  dioxide  gives 
up  oxygen  it  is  a  good  oxidizing  agent.  If  brought  in  contact 
with  substances  which  can  take  up  oxygen,  it  parts  with  it 
readily,  and  such  substances  are  oxidized.  Upon  this  fact  is 
based  its  value  as  a  disinfectant.  When  brought  in  contact  with 
organic  matter  it  oxidizes  i±  and  destroys  its  vitality.  Bacteria 
and  other  germs  which  produce  disease  are  thus  destroyed  by 
hydrogen  dioxide. 


CHAPTER  VIII 

COMPOUNDS    OP    CHLORINE     WITH     HYDROGEN    AND 

OXYGEN 

Hydrochloric  Acid,  HC1.  —  The  compound  formed  by  the 
union  of  hydrogen  with  chlorine  is  known  as  hydrochloric  acid. 
This  compound  is  formed  when  hydrogen  is  burned  in  chlorine. 
The  same  compound  is  formed  when  a  mixture  of  hydrogen 
and  chlorine  is  exposed  to  diffuse  light.  When  the  mixture  is 
exposed  to  direct  sunlight  they  combine  with  explosive  violence. 

Volume  Relations  in  which  Hydrogen  and  Chlorine  combine. 
—  It  will  be  remembered  that  one  volume  of  oxygen  combines 
with  two  volumes  of  hydrogen,  and  forms  two  volumes  of 
water-vapor.  The  relations  which  obtain  in  the  case  of  hydro- 
gen and  chlorine  are  even  simpler.  When  one  volume  of 
hydrogen  is  mixed  with  one  volume  of  chlorine  and  combina- 
tion takes  place,  all  of  both  gases  are  used  up,  and  just  two 
volumes  of  hydrochloric  acid  are  formed.  The  law  of  the 
simple  volume  relations  in  which  gases  combine  holds  here 
even  more  strikingly  than  in  the  case  of  oxygen  and  hydrogen. 

Preparation  of  Hydrochloric  Acid.  —  Hydrochloric  acid  gas 
is  prepared  most  conveniently,  as  follows:  When  a  salt  of 
hydrochloric  acid  is  treated  with  a  non-volatile  acid  such  as 
sulphuric,  hydrochloric  acid  gas  is  set  free.  The  best  known 
salt  of  hydrochloric  acid  is,  as  we  have  seen,  sodium  chloride. 
When  this  is  treated  with  sulphuric  acid,  a  reaction  takes 
place  in  the  sense  of  the  following  equation :  — 

2  NaCl  +  H2S04  =  Na2S04  +  2  HC1. 

The  hydrochloric  acid  gas  thus  formed  is  conducted  into 
water,  which  has  the  power  of  absorbing  large  (quantities  of 

62 


COMPOUNDS  OF  CHLORINE   WITH  HYDROGEN,  ETC.      63 

it.  This  is  the  form  in  which  it  is  used  in  the  laboratory  and 
in  the  arts. 

Chemical  Properties  of  Hydrochloric  Acid.  —  Hydrochloric 
acid,  as  the  name  implies,  is  an  acid,  and  since  this  is  the  first 
substance  which  we  have  thus  far  encountered  with  acid 
properties  (say  the  power  to  color  blue  litmus  red),  a  few  words 
should  be  added  in  reference  to  acids  in  general.  Hydrochloric 
acid  has  the  composition  represented  by  the  formula  HC1.  Its 
molecule  therefore  contains  one  atom  of  hydrogen  and  one  of 
chlorine.  The  question  arises  to  which  constituent  are  the 
acid  properties  due?  It  may  be  due  to  either  or  to  both. 
When  we  come  to  study  other  acids  we  shall  learn  that  many 
substances  are  acids  which  do  not  contain  chlorine,  and  many 
compounds  containing  chlorine  are  not  acids,  therefore,  chlorine 
is  not  essential  to  acidity.  We  shall  also  learn  that  all  sub- 
stances which  are  acid  contain  hydrogen,  and  no  other  element 
in  common.  Hydrogen  is  therefore  essential  to  acidity. 

There  are,  however,  many  compounds  which  contain  hydro- 
gen and  which  are  not  acids.  The  question  which  arises  is, 
How  does  the  hydrogen  in  the  latter  class  of  compounds  differ 
from  the  hydrogen  in  the  former  ? 

The  Theory  of  Electrolytic  Dissociation. — It  is  obvious  that 
the  atomic  theory  does  not  enable  us  to  answer  questions  like 
the  above.  In  terms  of  this  theory  one  atom  of  hydrogen  is 
like  any  other  atom  of  hydrogen,  and  the  chemical  behavior  of 
all  atoms  of  hydrogen  should  be  the  same.  The  fact  is  they 
are  very  different,  and  we  must,  therefore,  supplement  the 
atomic  theory  in  order  to  be  able  to  account  for  these 
differences. 

When  a  substance  like  hydrochloric  acid  is  dissolved  in 
water,  all  the  properties  of  the  solution  show  that  there  is  a 
larger  number  of  parts  in  the  solution  than  would  correspond 
to  the  molecules  of  the  compound.  It  would  lead  us  too  far 
in  this  connection  to  discuss  the  various  physical  chemical 
methods  by  which  this  conclusion  is  established.  Suffice  it  to 
say  here  that  this  conclusion  is  reached  as  the  result  of  the 


64  ELEMENTS  OF  INORGANIC   CHEMISTRY 

application  of  several  independent  methods,  and  is  established 
beyond  question. 

Since  a  solution  of  hydrochloric  acid  contains  a  larger 
number  of  parts  than  would  correspond  to  the  molecules  of 
the  compound,  it  is  obvious  that  the  molecules  in  solution  must 
be  broken  down  into  simpler  parts.  In  the  case  of  hydro- 
chloric acid  the  only  parts  into  which  the  molecule  could  break 
down  are  hydrogen  and  chlorine. 

Such  a  solution  is  capable  of  conducting  the  electric  current, 
and  the  only  way  in  which  a  solution  can  conduct  the  current 
is  for  the  dissolved  particles  to  carry  an  electrical  charge. 
Molecules  in  solution  do  not  conduct  the  current,  shown  by 
the  fact  that  solutions  of  cane  sugar,  alcohol,  and  the  like  are 
non-conductors.  '  The  parts  into  which  the  molecules  of  hydro- 
chloric acid  break  down  are,  therefore,  charged,  and  thus  differ 
from  ordinary  hydrogen  and  chlorine,  which  are  uncharged. 
These  charged  parts  are  known  as  ions.  The  ions  which  are 
charged  positively  are  known  as  cations,  while  those  which  are 
charged  negatively  are  known  as  anions.  When  a  molecule 
breaks  down  or  dissociates  as  we  say  in  the  presence  of  a 
solvent  like  water,  one  part  of  the  molecule  is  always  charged 
positively  or  is  a  cation,  while  the  other  part  is  always  charged 
negatively  or  is  an  anion.  This  theory  of  the  dissociation  of 
molecules  into  ions  is  known  as  the  theory  of  electrolytic  disso- 
ciation. 

Importance  of  the  Theory  of  Electrolytic  Dissociation  for 
Chemistry.  —  It  is  obvious  from  the  above  that  the  atomic 
theory  has  been  supplemented  by  a  new  conception.  In  addi- 
tion to  the  conception  of  the  atom  we  have  now  that  of  the 
atom  carrying  an  electrical  charge.  The  question  may  fairly 
be  asked,  why  bring  in  this  new  idea  ?  The  answer  is,  first, 
because  it  is  true.  We  have  an  overwhelming  amount  of  evi- 
dence in  favor  of  the  theory  of  electrolytic  dissociation,  most 
of  which,  obviously,  cannot  be  presented  here.  It  can,  how- 
ever, be  shown  that  molecules,  in  general,  cannot  react  chemi- 
cally. Dissolve  hydrochloric  acid  gas  in  some  solvent  like 


COMPOUNDS  OF  CHLORINE  WITH  HYDROGEN,  ETC.      65 

benzene,  which  does  not  have  the  property  to  any  marked 
degree  of  breaking  down  the  molecules  of  the  acid  into  ions. 
Such  a  solution  has  only  slightly  acid  properties.  It  does  not 
color  blue  litmus  red.  It  is  obvious  that  in  such  a  solution  we 
have  an  abundance  of  molecules  of  hydrochloric  acid,  and  yet 
such  slight  acid  properties.  The  molecules,  then,  do  not  have 
acid  properties. 

Again,  take  pure,  dry,  liquid  hydrochloric  acid.  We  have 
here  a  great  abundance  of  molecules  of  hydrochloric  acid,  and 
yet  no  acid  properties. 

On  the  other  hand,  wherever  we  have  the  molecules  of 
hydrochloric  acid  dissociated  into  ions  we  have  acid  properties 
manifesting  themselves  to  a  marked  extent. 

The  relations  pointed  out  for  hydrochloric  acid  hold  for  most 
chemical  substances.  The  molecules  are  relatively  inert;  the 
ions  being  the  active  agents  chemically.  It  is  thus  obvious  that 
the  theory  of  electrolytic  dissociation  is  of  fundamental  im- 
portance for  the  whole  science  of  chemistry. 

Method  of  Indicating  the  Existence  of  Ions. — As  we  shall 
have  to  deal  frequently  with  ions,  we  adopt  some  method  of 
distinguishing  between  atoms  and  ions.  Since  ions  are  charged 
atoms  or  groups  of  atoms,  we  shall  use  the  positive  sign  over 
the  symbol  of  an  atom  to  mean  that  it  is  charged  positively, 
or  is  a  cation.  The  negative  sign  over  an  atom  or  group  of 
atoms  means  that  it  is  carrying  a  negative  charge  and  is  an 
anion.  Hydrochloric  acid  is  dissociated  by  water  in  the  sense 
of  the  following  equation :  — 

HC1  =  H,  Cl. 

The  comma  between  the  two  iong  means  that  they  come  from 
the  same  molecule,  and  will  unite  again  and  form  the  original 
molecule  when  the  water  is  removed.  This  method  of  express- 
ing the  dissociation  of  molecules  with  ions  is  adopted  through- 
out the  book. 

Action  of  Hydrochloric  Acid  on  a  Metal  like  Zinc.  —  When  a 
solution  of  hydrochloric  acid  is  brought  in  contact  with  a  metal 


66  ELEMENTS  OF  INORGANIC  CHEMISTRY 

like  zinc,  the  latter  takes  the  positive  charge  from  the  hydrogen 
ion,  becoming  itself  an  ion  and  passing  into  solution,  while 
the  hydrogen  ion  having  lost  its  electrical  charge  becomes  an 
atom.  Two  atoms  of  hydrogen  then  combine  and  form  a  mole- 
cule. The  reaction  between  zinc  and  hydrochloric  acid  is  rep- 
resented by  the  following  equation  :  — 


Zn  +  H,  Cl  +  H,  01  =  Zn,  01,  C1  +  H 


Hydrochloric  acid  acts  upon  metals  in  general  in  the  sense 
of  the  above  equation  —  the  metal  taking  the  charge  from  the 
hydrogen  ion,  becoming  itself  an  ion,  converting  the  hydrogen 
into  the  atomic  condition,  and  this  then  passes  into  the  molec- 
ular condition. 

Definition  of  an  Acid.  —  Having  studied  hydrochloric  acid  as 
the  type  of  a  large  class  of  chemical  compounds  known  as 
acids,  we  are  prepared  to  consider  these  a  little  more  closely. 

All  acids  have  certain  properties  in  common.  They  all  taste 
sour  ;  they  have  the  property  of  coloring  certain  vegetable  dyes 
red  ;  they  all  contain  hydrogen  which  can  give  up  its  electri- 
cal charge  to  certain  metals,  itself  escaping  as  hydrogen  gas. 

Hydrogen  in  this  form  is  known  as  ionic  hydrogen,  and,  as 
has  been  stated,  wherever  we  have  ionic  hydrogen  we  have 
acid  properties,  and  wherever  we  have  acid  properties  we 
have  ionic  hydrogen.  To  say  that  a  compound  has  acid  prop- 
erties means,  then,  that  when  it  is  dissolved  in  ivater  or  some 
other  dissociating  solvent  it  yields  hydrogen  ions. 

This  definition  says  that  a  compound  is  not  an  acid  unless  it 
is  brought  into  the  presence  of  a  dissociating  solvent.  This  is 
the  same  as  to  say  that  no  pure,  homogeneous  substance  is  an 
acid. 

We  shall  learn  that  the  above  relations  hold  for  all  acids. 
No  pure,  dry,  homogeneous  substance  is  an  acid.  It  becomes 
an  acid  only  when  dissociated  by  a  solvent  or  some  other 
means  into  hydrogen  cations,  and  into  anions  whose  nature 
depends  upon  the  acid  in  question  and  varies  with  every  acid. 


COMPOUNDS  OF  CHLORINE  WITH  HYDROGEN,  ETC.      67 

Detection  of  Hydrochloric  Acid.  —  There  is  one  reaction 
which  serves  to  detect  hydrochloric  acid  under  all  ordinary 
conditions.  Hydrochloric  acid  is,  as  we  have  seen,  dissociated 
into  hydrogen  ions  and  chlorine  ions. 

To  detect  hydrochloric  acid  we  must  make  use  of  some  re- 
action which  is  characteristic  of  the  chlorine  ion,  since  all  acids 
yield  hydrogen  ions.  Such  a  reaction  takes  place  whenever  a 
silver  ion  is  brought  in  contact  with  a  chlorine  ion. 

H,  Cl  +  Ag,  N03  =  AgCl  +  H,  N03 

expresses  the  reaction  between  hydrochloric  acid  and  silver 
nitrate.  The  silver  chloride  formed  is  a  white  solid,  readily 
soluble  in  ammonia.  Since  soluble  chlorides  in  general  yield 
in  solution  chlorine  ions,  this  is  a  means  of  detecting  also  the 
presence  of  such  chlorides. 

Physical  Properties  of  Hydrochloric  Acid.  —  Hydrochloric 
acid  is  a  colorless  gas,  with  a  sharp,  pungent  odor,  and  pro- 
duces marked  irritation  of  the  mucous  membrane  of  the  nose 
and  throat  when  inhaled  even  in  small  quantity.  The  gas 
shows  unusual  solubility  in  water,  one  -volume  of  water  at 
zero  degrees  dissolving  about  503  volumes  of  the  gas.  The 
gas  has  such  great  attraction  for  water  that  if  the  breath 
is  blown  across  the  open  mouth  of  a  bottle  of  concentrated 
hydrochloric  acid,  the  particles  of  water  are  condensed  by  the 
escaping  hydrochloric  acid  and  a  mist  is  produced,  which  can 
be  readily  seen.  The  same  effect  is  observed  when  the  breath 
is  blown  into  a  stream  of  hydrochloric  acid  gas  escaping  from 
a  generator. 

Aqueous  Solution  of  Hydrochloric  Acid.  —  The  hydrochloric 
acid  that  is  used  in  the  laboratory  is  a  solution  of  the  gas  in 
water.  The  concentration  of  the  solution  increases  as  more  and 
more  gas  is  dissolved. 

Compounds  of  Chlorine  with  Oxygen.  —  Although  chlorine 
and  oxygen  cannot  be  made  to  combine  directly,  several  com- 
pounds of  these  two  elements  have  been  made  by  indirect 


68  ELEMENTS  OF   INORGANIC   CHEMISTRY 

methods.  These  compounds  are  chlorine  monoxide  (C120), 
chlorine  dioxide  (C102),  and  chlorine  septoxide  (C1207).  These 
compounds  are  all  characterized  by  instability. 

Compounds  of  Chlorine  with  Hydrogen  and  Oxygen.  —  Thus 
far  we  have  studied  compounds  between  only  two  elements. 
We  might  suspect,  therefore^  that  only  two  chemical  elements 
can  combine  with  one  another,  forming  a  definite  molecule,  or 
at  least  that  this  is  by  far  the  most  common  form  of  chemical 
union.  Such  is  by  no  means  the  case.  We  shall  now  study 
briefly  a  class  of  compounds  between  the  three  elements, 
chlorine,  oxygen,  and  hydrogen,  which,  if  not  very  common 
substances,  have  considerable  chemical  interest.  These  com- 
pounds are :  — 

Hypochlorous  acid HC10. 

Chlorous  acid HC102. 

Chloric  acid HC103. 

Perchloric  acid HC104. 

Hypochlorous  Acid,  HOC1.  —  When  chlorine  is  conducted  into 
a  cold,  dilute  solution  of  potassium  hydroxide,  the  reaction 
takes  place  in  the  sense  of  the  following  equation:  — 

2  KOH  +  C12  =  KC1  +  KOC1  +  H20. 

The  salts  of  hypochlorous  acid  are  termed  hypochlorites.  This 
salt  is,  therefore,  potassium  hypochlorite. 

When  this  salt  is  treated  with  a  cold,  dilute  solution  of 
hydrochloric  acid,  the  hypochlorous  acid  is  set  free:  — 

KOC1 4-  HC1  =  KC1  +  HOC1. 

The  hypochlorous  acid  can  then  be  distilled  off  and  collected. 

Calcium  Hypochlorite,  Ca(OCl)2. — The  calcium  salt  of  hypo- 
chlorous  acid  is  used  extensively  as  a  bleaching  agent  on  account 
of  the  ease  with  which  it  gives  up  chlorine.  When  chlorine  is 
passed  into  slaked  lime,  the  following  reaction  apparently 
takes  place :  — 

2  Ca(OH)2  +  2  C12  =  CaCl2  +  Ca(OCl)2  +  2  H20. 


COMPOUNDS  OF  CHLORINE  WITH  HYDROGEN,  ETC.      69 

This  apparent  mixture  of  calcium  chloride  and  calcium  hypo- 
chlorite  is  known  as  bleaching-powder,  and  is  largely  used  as 
a  disinfectant. 

Chloric  Acid,  HC103. — When  chlorine  gas  is  conducted  into  a 
hot,  concentrated  solution  of  potassium  hydroxide,  the  follow- 
ing reaction  takes  place :  — 

6  KOH  +  3  C12  =  5  KC1  +  KC103  +  3  H20. 

The  solution  contains,  after  the  reaction  is  over,  two  salts, 
potassium  chloride  and  potassium  chlorate.  These  can,  how- 
ever, be  readily  separated  by  their  different  solubilities  in 
water  —  potassium  chloride  being  quite  soluble,  while  potas- 
sium chlorate  is  very  much  less  soluble. 

From  the  solution  potassium  chlorate  readily  crystallizes, 
especially  on  evaporation,  leaving  behind  in  solution  the  potas- 
sium chloride.  With  potassium  chlorate  we  have  already  become 
somewhat  familiar  when  we  were  studying  methods  of  prepar- 
ing oxygen.  It  will  be  remembered  that  this  compound  gives 
off  all  of  its  oxygen  when  heated  to  an  elevated  temperature. 

When  potassium  chlorate  is  treated  with  a  dilute  solution  of 
sulphuric  acid,  the  following  reactioii  takes  place :  — 

2  KC103  +  H2S04  =  K2S04  +  2  HC103. 

Care  must  be  taken  not  to  treat  potassium  chlorate  with  concen- 
trated sulphuric  acid,  since  violent  explosions  almost  always  result. 

The  solution  contains  the  chloric  acid,  but  since  the  latter 
cannot  be  distilled  without  undergoing  decomposition,  this 
method  does  not  yield  pure  chloric  acid. 

To  obtain  pure  chloric  acid  a  salt  of  this  acid  must  be 
used  which  will  form  an  insoluble  precipitate  with  sulphuric 
acid.  The  barium  salt  is  the  most  convenient.  When  barium 
chlorate  is  treated  with  a  dilute  solution  of  sulphuric  acid  in 
equivalent  quantity,  insoluble  barium  sulphate  is  precipitated, 
and  pure  chloric  acid  remains  in  solution. 

Ba(C103)2  +  H2S04  =  BaS04  +  2  HC103. 

(Insoluble) 


70  ELEMENTS  OF  INORGANIC  CHEMISTRY 

The  barium  sulphate  is  then  filtered  off,  or  the  clear,  superna- 
tant liquid  decanted  from  the  precipitate  and  concentrated  in  a 
vacuum  or  over  sulphuric  acid. 

Properties  of  Chloric  Acid.  —  Chloric  acid  is  a  colorless  liquid, 
with  very  strongly  acid  properties  and  with  great  oxidizing 
power.  It  contains  a  large  amount  of  oxygen,  which  it  readily 
gives  up.  When  a  piece  of  paper  is  saturated  with  a  concen- 
trated solution  of  chloric  acid,  it  is  oxidized  so  energetically 
that  it  bursts  into  flame. 

Chlorates.  —  The  chlorates,  as  the  salts  of  chloric  acid  are 
called,  are  very  energetic  oxidizing  agents.  This  is  due  in 
part  to  the  large  amount  of  oxygen  which  they  contain,  and  in 
part  to  the  ease  with  which  they  give  it  up. 

The  Chlorine  Ion  and  the  Ion  of  Chlorates.  —  We  have  studied 
one  of  the  characteristic  properties  of  the  chlorine  ion,  viz.  its 
power  to  combine  with  the  silver  ion  and  form  insoluble  silver 
chloride.  We  would  naturally  ask  whether  the  chlorine  in 
potassium  chlorate  has  this  same  power.  Potassium  chlorate 
dissociates  as  follows:- — 

KC103  =  K,  Cf03. 

Chlorine  in  this  case,  instead  of  forming  th'e  anion,  forms  only 
a  part  of  the  anion.  It  is  in  combination  with  three  oxygen 
atoms,  and  the  chlorine  and  oxygen  form  the  anion.  If  a  solu- 
tion of  potassium  chlorate  is  treated  with  a  solution  of  silver 
nitrate,  no  precipitate  is  formed;  showing  that  chlorine  has 
very  different  properties  when  alone  in  the  ionic  state,  than 
when  combined  with  another  element  forming  part  of  a  com- 
plex ion. 

Perchloric  Acid,  HC104.  —  When  potassium  chlorate  is  heated 
vigorously,  it  gives  off  all  of  its  oxygen,  as  we  saw  when  we 
were  studying  methods  for  the  preparation  of  oxygen.  If, 
however,  potassium  chlorate  is  heated  moderately,  it  gives  off 
only  a  part  of  its  oxygen. 

2  KC103  =  KC1  +  KC104  +  02. 


COMPOUNDS  OF  CHLORINE  WITH   HYDROGEN,  ETC.      71 

The  perchlorate,  KC104,  can  be  obtained  from  the  mixture  by 
dissolving  out  the  potassium  chloride  by  means  of  cold  water 
in  which  potassium  chlorate  is  very  slightly  soluble. 

When  potassium  perchlorate  is  treated  with  sulphuric  acid, 
perchloric  acid  is  set  free. 

2  KC104  +  H,S04  =  K2S04  +  2  HC104. 


Properties  of  Perchloric  Acid.  —  Perchloric  acid  is  more 
stable  than  any  of  the  other  oxygen  and  hydrogen  compounds 
of  chlorine.  The  pure  acid  is  a  very  vigorous  oxidizing  agent, 
and  explodes  easily  when  brought  in  contact  with  substances 
which  can  be  oxidized. 

Perchloric  acid  is  a  very  strong  acid,  which  is  the  same  as 
to  say  that  it  is  very  much  dissociated  by  water.  It  readily 
replaces  hydrochloric  acid  from  its  salts,  but  this  is  partly  due 
to  the  comparative  insolubility  of  the  perchlorates  ;  it  being  a 
general  law  in  chemistry  that  when  an  insoluble  compound 
can  be  formed  it  is  formed.  If  a  fairly  concentrated  solution  of 
potassium  chloride  is  treated  with  perchloric  acid,  needles  of 
potassium  perchlorate  are  precipitated. 

Valence.  —  Two  atoms  of  chlorine  are  sometimes  required  to 
combine  with  one  atom  of  oxygen,  as  in  the  compound  C120. 
In  other  cases,  one  atom  of  chlorine  combines  with  two  atoms 
of  oxygen,  as  in  C102,  while  in  the  compound  C1207,  one  atom 
of  chlorine  combines  with  three  and  a  half  atoms  of  oxygen. 
The  power  of  an  atom  to  hold  other  atoms  in  combination  is 
known  as  its  valence.  This  is  not  to  be  taken  as  a  definition, 
but  simply  as  a  description  of  the  action  of  chemical  valence. 

We  have  already  seen  examples  of  chemical  action  taking 
place  between  ions,  which  are  atoms  or  groups  of  atoms  charged 
with  electricity.  There  is  abundant  evidence  that  nearly  all 
chemical  action  is  between  ions  or  charged  parts.  The  estab- 
lishment of  this  fact  has  given  us  a  definite  physical  basis  for 
the  conception  of  valence. 

An  ion  is  univalent,  or  can  combine  with  one  hydrogen  atom, 
which  is  really  the  unit  of  valence,  if  it  carries  one  electrical 


72  ELEMENTS   OF   INORGANIC   CHEMISTRY 

charge.  An  ion  is  bivalent,  or  can  combine  with  two  hydrogen 
atoms,  which  are  equivalent  under  all  ordinary  conditions  to 
one  oxygen  atom,  if  it  carries  two  electrical  charges ;  trivalent, 
if  it  carries  three  charges;  quadrivalent,  if  it  carries  four 
charges;  quinquivalent,  if  there  are  five  charges  upon  it;  sex- 
ivalent,  if  it  carries  six  charges ;  septivalent,  if  there  are  seven 
charges  connected  with  it;  and  octivalent,  if  it  carries  eight 
charges.  There  are  no  ions  known  which  have  a  greater  val- 
ence than  eight;  i.e.  which  have  the  power  of  combining  with 
more  than  eight  hydrogen  atoms,  or  four  oxygen  atoms. 

Faraday's  Law  the  Basis  of  Chemical  Valence.  —  Faraday 
passed  an  electric  current  through  a  solution  of  an  electrolyte, 
and  observed  that  the  amount  of  the  electrolyte  decomposed  was 
proportional  to  the  amount  of  current  which  had  passed  through 
the  solution.  On  the  basis  of  this  experimentally  established 
fact  he  enunciated  the  first  part  of  his  well-known  law  :  — 

TJie  amount  of  chemical  decomposition  effected  by  the  passage 
of  the  current  is  proportional  to  the  amount  of  electricity  which 
flows  through  the  conductor. 

Since  electricity  can  flow  through  a  solution  of  an  electrolyte 
only  by  being  carried  by  the  ions  in  the  solution,  the  above  part 
of  Faraday's  law  shows  that  each  ion  of  the  same  substance 
carries  exactly  the  same  amount  of  electrical  energy. 

Faraday  determined  also  the  amounts  of  different  elements 
which  are  separated  from  their  compounds,  by  passing  the 
same  current  through  solutions  of  these  compounds.  A  gen- 
eralization of  very  wide  significance  was  reached,  which  is  the 
second  part  of  the  law  of  Faraday  :  — 

TJie  amounts  of  the  different  elements  which  are  separated  by 
the  same  quantity  of  electricity  bear  the  same  relation  to  one  an- 
other as  the  equivalents  of  these  elements. 

This  is  saying  in  other  words  that  all  univalent  elements 
carry  exactly  the  same  quantity  of  electricity,  all  bivalent  ele- 
ments carry  exactly  twice  this  quantity,  all  trivalent  elements 


COMPOUNDS  OF  CHLORINE  WITH  HYDROGEN,  ETC.      73 

three  times  the  quantity,  and  so  on.  In  a  word,  all  univalent 
ions  carry  the  same  amount  of  electricity,  and  all  polyvalent 
ions  a  simple,  rational  multiple  of  the  amount  carried  by  univ- 
alent ions  —  the  multiple  being  the  valence  of  the  ion. 


EXPERIMENTS  WITH  THE  COMPOUNDS  OF  CHLORINE 
WITH  HYDROGEN  AND  OXYGEN 

Experiment  32.     Preparation  of  Hydrochloric  Acid  by  the 
Action  of  Sulphuric  Acid  on  Sodium  Chloride.  —  (Same  apparatus 
as  in  Experiment  26 ;  sodium  chloride ;  sulphuric  acid  contain- 
ing by  volume  2  parts  of 
acid  to  1  of  water.) 

Arrange  an  apparatus 
as  in  Experiment  26.  In- 
troduce into  the  flask  A 
60  grams  of  sodium  chlo- 
ride. Pour  in  gradually 
and  cautiously  about  75 
grams  of  the  acid  of  the 
concentration  given  above. 
Warm  gently.  Hydro- 
chloric acid  will  be  given 
off  in  abundance.  Vessels 
can  be  filled  with  hydro- 
chloric acid  in  the  same 
manner  that  they  were 
filled  with  chlorine,  since 
the  gas  is  heavier  than  air. 
Fill  a  number  of  cylinders 
and  bottles  with  the  gas 
for  further  experiments. 
(See  page  62.) 

Experiment  33.  Prep- 
aration of  Hydrochloric 
Acid  Gas  by  the  Action 
of  Concentrated  Sulphuric 
Acid  on  a  Concentrated 
Aqueous  Solution  of  Hydrochloric  Acid.  —  (Dropping-funnel ; 
500  cc.  flask ;  exit-tube ;  rubber  stopper ;  small  funnel ;  con- 
centrated hydrochloric  acid;  concentrated  sulphuric  acid.) 

Arrange  the  apparatus  as  shown  in  Figure  23.     Allow  the 


FIG.  23. 


74  ELEMENTS   OF  INORGANIC   CHEMISTRY 

sulphuric  acid  to  drop  very  slowly  into  the  concentrated  hydro- 
chloric acid.  Collect  the  gas  in  water,  allowing  the  funnel  at 
the  end  of  the  exit-tube  to  dip  just  beneath  the  surface  of  the 
water.  Preserve  the  aqueous  solution  of  hydrochloric  acid 
for  further  experiments. 

Place  a  thermometer  in  the  water  before  the  gas  is  conducted 
in,  and  observe  the  rise  in  temperature  as  the  gas  is  absorbed. 

Experiment  34.  Properties  of  an  Aqueous  Solution  of  Hydro- 
chloric Acid.  —  (Aqueous  solution  of  hydrochloric  acid ;  potas- 
sium, hydroxide;  solution  of  litmus;  alcoholic  solution  of 
phenolphthalein ;  piece  of  zinc ;  solution  of  silver  nitrate.) 

Add  some  of  the  hydrochloric  acid  to  a  piece  of  zinc.  Show 
that  hydrogen  escapes. 

Add  a  drop  of  the  solution  of  hydrochloric  acid  to  silver 
nitrate.  Observe  the  white  precipitate  formed.  Add  ammonia, 
and  the  white  precipitate  will  dissolve. 

Add  a  few  drops  of  the  solution  of  hydrochloric  acid  to 
much  water.  Taste  the  solution.  Add  a  few  drops  of  the 
solution  of  litmus  and  observe  the  red  color.  Add  drop  by 
drop  the  solution  of  potassium  hydroxide,  stirring  the  solution 
thoroughly  after  the  addition  of  each  drop  until  the  solution 
turns  blue.  To  the  blue  solution  add  a  little  more  hydro- 
chloric acid  and  observe  the  red  color  appear  again. 

To  another  portion  of  very  dilute  hydrochloric  acid  add  a 
few  drops  of  the  alcoholic  solution  of  phenolphthalein.  The 
solution  remains  colorless.  Add  drop  by  drop  potassium 
hydroxide.  The  solution  becomes  beautifully  red.  Add  more 
of  the  acid,  and  the  red  color  will  disappear.  These  are  some 
of  the  reactions  of  hydrochloric  acid :  Action  on  metals. 
Formation  of  an  insoluble  compound  with  silver  nitrate. 
Neutralization  of  bases.  (See  page  67.) 

Experiment  35.  Preparation  of  Calcium  Hypochlorite ; 
Bleaching-powder.  —  (Flask  holding  250  cc.  lime;  chlorine 
generator.) 

Add  water  drop  by  drop  to  25  grams  of  lime  until  it  is  just 
slaked,  but  is  still  a  dry  powder.  Introduce  into  a  flask 
holding  250  cc.  and  pass  chlorine  into  the  flask,  shaking  the 
dry  slaked  lime  very  frequently. 

Bleaching-powder  is  formed  in  a  short  time.  Treat  a  little 
with  hydrochloric  acid.  What  is  given  off  ?  (See  page  68.) 

Experiment  36.  Preparation  of  Potassium  Chlorate.  — 
(Chlorine  generator ;  beaker ;  caustic  potash.) 

Dissolve  50  grams  of  caustic  potash  in  100  cc.  of  water  and 


COMPOUNDS  OF  CHLORINE  WITH  HYDROGEN,  ETC.      75 

pass  chlorine  into  the  solution  as  long  as  it  is  absorbed. 
Boil  off  about  one-half  of  the  liquid.  Crystals  of  potassium 
chlorate  are  deposited  from  the  solution  when  cold,  together 
with  some  potassium  chloride.  Filter  off  the  crystals  and  dis- 
solve them  in  a  minimum  amount  of  hot  water.  Allow  the 
solution  to  cool,  when  practically  pure  potassium  chlorate  will 
crystallize  out. 

Most  of  the  potassium  chloride  formed  remains  in  the  original 
solution,  since  potassium  chloride  is  so  readily  soluble  in  water. 
Evaporate  the  solution  to  about  one-fourth  volume,  to  allow 
most  of  the  potassium  chlorate  remaining  in  solution  to  separate. 
Filter  off  this  substance  and  evaporate  the  remaining  solution 
to  dryness.  Potassium  chloride  will  separate  from  the  solu- 
tion. 

Test  the  potassium  chlorate  by  heating  a  small  portion  in  a 
hard-glass  tube,  and  show  that  oxygen  is  given  off. 

Test  the  potassium  chloride  by  treating  a  little  in  a  test-tube 
with  concentrated  sulphuric  acid,  and  show  that  hydrochloric 
acid  escapes.  (See  page  69.) 

Experiment  37.  Oxidizing  Power  of  Chloric  Acid.  —  (Pipette ; 
potassium  chlorate ;  cane  sugar ;  concentrated  sulphuric  acid.) 

( This  experiment  to  be  performed  only  by  the  instructor.)  Mix. 
a  few  grams  of  cane  sugar  with  a  few  grams  of  potassium 
chlorate,  and  place  the  mixture  on  a  stone  slab  under  the  hood. 
Add  cautiously  from  the  pipette  a  drop  or  two  of  concentrated 
sulphuric  acid.  The  whole  mass  will  burst  into  violent  flame. 
(See  page  70.) 


CHAPTER   IX 

NITROGEN 

*  Occurrence  and  Preparation.  —  The  chief  source  of  nitrogen 
is  the  atmospheric  air,  which  consists  approximately  of  one-fifth 
oxygen  and  four-fifths  nitrogen. 

It  can  be  prepared  in  fairly  pure  condition  by  removing  the 
oxygen  from  atmospheric  air.  This  can  be  accomplished  by 
means  of  phosphorus.  When  moist  phosphorus  is  brought  in 
contact  with  the  air,  the  oxygen  combines  with  the  phosphorus, 
and  the  nitrogen  remains  behind. 

The  oxygen  can  be  removed  from  atmospheric  air  also  by  cer- 
tain metals  at  an  elevated  temperature.  Thus,  when  metallic 
copper  is  heated  to  redness  in  the  presence  of  atmospheric  air, 
the  oxygen  combines  with  the  copper,  forming  cupric  oxide, 
CuO,  and  the  nitrogen  remains  behind.  ^ 

To  prepare  nitrogen  of  a  high  degree  of  purity,  certain  chemi- 
cal reactions  are  made  use  of.  When  ammonium  nitrite,  a 
compound  having  the  composition  NH4N02,  is  heated,  the  fol- 
lowing reaction  takes  place :  — 

NH4N02==2H20-f]Sr2. 

This  is  an  excellent  method  for  preparing  pure  nitrogen. 
X  Properties  of  Nitrogen.  —  Nitrogen  is  characterized  by  its 
inertness,  not  only  at  ordinary  temperatures,  but  even  at  elevated 
temperatures.     On  account  of  its  chemical  inactivity  nitrogen 
cannot  support  combustion. 

It  cannot  support  life,  all  animals  dying  in  a  very  short  time 
when  compelled  to  breathe  only  nitrogen.  Nitrogen  is  a  taste- 
less, odorless,  colorless  gas.;*  It  forms  a  colorless  liquid,  boiling 
at  — 196°.  One  litre  of  nitrogen  weighs  1.2521  grams.  When 
liquid  nitrogen  is  cooled  to  —  214°,  it  solidifies. 

76 


NITROGEN  77 


COMPOUND  OF  NITROGEN  WITH  HYDROGEN 

Ammonia,  NH3.  —  The  best-known  compound  of  nitrogen  and 
hydrogen  is  ammonia.  Ammonia  occurs  in  nature  in  small 
quantities  in  the  free  condition.  In  the  form  of  its  salts  it 
occurs  in  many  soils,  and  on  account  of  their  great  solubility 
the  salts  of  ammonia  exist  largely  in  solution  in  water.  The 
ammonia  is  liberated  in  considerable  quantity  by  decomposing 
organic  matter.  This  is  readily  detected  by  the  odor  of  the 
gas  escaping  from  decomposing  animal  remains  or  decaying 
vegetable  matter. 

Ammonia  can  be  formed  by  the  direct  union  of  hydrogen  and 
nitrogen.  When  one  part  of  nitrogen  is  mixed  with  three  parts 
of  hydrogen  and  electric  sparks  passed  through  the  mixture,  a 
part  of  the  hydrogen  and  nitrogen  combine,  forming  ammonia:  — 

"VT     _j_  O  TT     O  "VTTJT 

1>  2  ~T  <5  ±1-2  =  •"  -^  -tJ-3- 

The  volume  of  the  ammonia  formed  is  just  half  the  sum  of  the 
volumes  of  the  nitrogen  and  hydrogen  which  have  disappeared. 
When  one  volume  of  nitrogen  combines  with  three  volumes  of 
hydrogen,  there  are  two  volumes  of  ammonia  formed. /*» 

This  shows  again  the  simple  relations  by  volume  in  which 
gases  combine,  and  the  simple  relation  between  the  volum*es  of 
the  original  gases  and  the  volume  of  the  product. 

Ammonia  is  most  conveniently  prepared  by  the  action  of  a 
base  on  an  ammonium  salt.  When  ammonium  chloride  is  boiled 
with  an  aqueous  solution  of  a  strong  base  like  sodium  hydroxide, 
ammonia  gas  is  liberated :  — 

NH4C1  +  NaOH  =  NaCl  +  H20  +  NH3. 

In  the  laboratory  ammonia  is  prepared  most  conveniently  by 
mixing  slaked  lime  with  ammonium  chloride,  and  warming  the 
mixture.  The  reaction  is  —  f 

1  +  Ca(OH)2  =  CaCl2  +  2  H20 

77 


78  ELEMENTS  OF   INORGANIC   CHEMISTRY 

Ammonia  was  formerly  obtained  from  decaying  organic  mat- 
ter, and  from  ammonium  salts  which  occurred  in  certain  arid 
regions  of  the  earth.  The  ammonium  chloride  which  occurred 
in  the  neighborhood  of  the  temple  of  Jupiter  Ammon  was 
termed  sal  ammoniac,  whence  the  origin  of  the  name  ammonia. 
/•/-Ammonia  to-day  is  obtained  mainly  by  the  dry  distillation  of 
coal  in  the  manufacture  of  illuminating  gas.>C 
^  Chemical  Properties  of  Ammonia.  —  Ammonia  in  the  pure,  dry 
condition  is  not  active  chemically.  When  perfectly  dry  ammonia 
gas  is  brought  in  contact  with  perfectly  dry  hydrochloric  acid  gas, 
there  is  not  the  slightest  reaction  between  the  two  substances.  If 
there  is  a  mere  trace  of  moisture  present,  the  two  gases  react 
at  once,  forming  the  solid  ammonium  chloride. 

Ammonia  dissolves  in  wate^1  with  the  greatest  ease,  forming 
a  compound  which  neutralizes  acids  and  which  is,  therefore,  a 
basic  substance.  X 

When  ammonia  dissolves  in  water,  it  combines  with  the  water, 
forming  the  compound  NH4OH  :  — 


The  compound  NH4OH  which,  however,  has  never  been  isolated, 
is  then  acted  on  by  more  water,  and  dissociates  thus  :  — 

NH4OH=NH4,OH. 

The  hypothetical  group  NH4  is  called  ammonium.  While  this 
group  has  not  been  isolated,  there  is  little  doubt  as  to  its  exist- 
ence. It  plays  the  same  role,  as  we  shall  see,  that  a  metal  atom 
does  in  the  formation  of  compounds. 

Physical  Properties  of  Ammonia.  —  Ammonia  is  a  colorless  gas 
with  a  very  penjeJrj,Ung_qdor.  One  litre  of  ammonia  at  0°  and 
760  mm.  pressure  weighs  0.775  grams. 

Liquid  ammonia  has  a  very  high  specific  heat,  which  is  the 
same  as  to  say  that  a  large  amount  of  heat  is  required  to  produce 
a  small  rise  in  temperature.  According  to  some  authorities, 
slightly  higher  even  than  water.  J*0n  account  of  its  high  specific 
heat  and  high  heat  of  vaporization  it  is  an  excellent  refriger- 


NITROGEN  79 

ating  agent,  and  is  used  extensively  for  this  purpose,  especially 
in  connection  with  the  artificial  preparation  of  ice&i  Ammonia 
gas  is  liquefied  by  pressure,  a  large  amount  of  heat  being, 
of  course,  set  free  during  the  process.  This  heat  is  removed 
by  a  current  of  cold  water  flowing  around  the  vessel  in  which 
the  liquefaction  is  taking  place.  The  liquid  ammonia  then 
flows  into  tubes  which  closely  surround  the  vessels  containing 
the  water  which  is  to  be  frozen,  and  is  allowed  to  vaporize  in 
these  tubes.  In  vaporizing  it  must  obtain  heat  from  some- 
where, and  takes  it  from  surrounding  objects.  The  water  loses 
its  heat,  is  cooled  below  the  freezing  temperature  and  solidifies. 
The  ammonia,  having  passed  into  the  form  of  vapor,  is  not 
lost,  but  is  pumped  into  the  liquefying  chamber,  subjected 
again  to  pressure  and  liquefied,  the  heat  set  free  being  again 
removed  by  the  current  of  cold  water.  The  process  is  thus  a 
continuous  one,  the  same  ammonia  being  used  over  and  over 
again. 

Machines  for  freezing  water  by  means  of  liquid  ammonia 
were  early  devised  by  Carre  and  are  known  as  Carre  ice 
machines.  Many  of  the  modern  devices  are  modifications  of 
these  machines  of  Carre,  utilizing  exactly  the  same  principles. 

A  few  years  ago  most  of  the  "  artificial  ice  "  was  made  by 
the  ammonia  process.  Now,  considerable  ice  is  obtained  by 
allowing  water  to  evaporate  into  a  space  under  diminished 
pressure. 

Ammonium  Amalgam,  NH4Hg.  —  While  the  compound  NH4 
has  never  been  isolated,  its  amalgam,  or  compound  with  mer- 
cury, is  readily  prepared.  When  sodium  amalgam,  a  com- 
pound of  sodium  and  mercury  having  the  composition  NaHg, 
is  treated  with  a  concentrated  solution  of  ammonium  chloride, 
the  amalgam  swells  up,  occupying  a  relatively  large  volume. 
The  product  has  a  metallic  lustre,  and  is  probably  ammonium 
amalgam.  The  reaction  probably  takes  place  in  the  sense  of 
the  following  equation  :  — 

NaHg  +  NH4C1  =  NaCl  +  NHJIg. 


80  ELEMENTS  OF   INORGANIC   CHEMISTRY 

It  would  seeni  that  ammonium  amalgam  was  a  hopeful  sub- 
stance from  which  to  obtain  the  group  ammonium.  It  is,  how- 
ever, unstable,  breaking  down  at  ordinary  temperatures  into 
ammonia,  hydrogen,  and  mercury. 

COMPOUNDS  OF  NITROGEN  WITH  OXYGEN  AND 
HYDROGEN 

Ammonium  Hydroxide,  NH4OH.  —  We  have  already  seen  that 
ammonium  hydroxide  is  a  basic  substance.  This  is  a  com- 
pound of  nitrogen  with  oxygen  and  hydrogen,  and  must  be 
considered  here.  The  most  characteristic  property  of  this  sub- 
stance is  its  basic  nature.  Being  a  base,  it  neutralizes  acids 
forming  salts.  Ammonium  hydroxide  unites  readily  with 
hydrochloric  acid,  forming  the  well-characterized,  beautifully 
white  salt,  ammonium .  chloride :  — 

NH4,  0~H  +  H,  Cl  =  H20  +  NH4,  Cl. 

• 

As  the  water  is  removed,  the  ammonium  and  chlorine  ions 
combine,  forming  ammonium  chloride  :  — 

NH4,  Cl  =  NH4C1. 

*kCompounds  of  Nitrogen  with  Oxygen.  —  Nitrogen  forms  the 
following  compounds  with  oxygen  :  Nitrous  oxide,  N20 ;  nitric 
ojdde,  NO ;   nitrogen  sesquioxide  or  trioxide,  N203 ;    nitrogen 
djoxide  or  tetroxide,  depending  upon  whether  it  has  the  com- 
position N02  or  N204 ;  and  nitrogen  pentoxide,  N205xf 
We  shall  now  study  these  compounds  in  some  detaiL 
Nitrous  Oxide,  N20. — Nitrous  oxide  is  formed  when  am- 
monium nitrate  is  heated  to  250°.     The  following  equation 
expresses  the  reaction  that  takes  place :  — 

NH4N03  =  2  H20  +  N20. 

The  oxygen  and  hydrogen  combine  and  form  water,  and  the 
nitrogen  and  oxygen  form  the  compound  N20,  which  escapes. 


NITROGEN  81 

Nitrous  oxide  is  a  remarkable  substance,  in  that  it  supports 
combustion  almost  as  well  as  pure  oxygen.  Phosphorus  and 
carbon  burn  readily  in  nitrous  oxide.  Certain  substances,  how- 
ever, burn  in  oxygen  and  burn,  but  not  so  readily,  in  nitrous 
oxide.  The  products  of  combustion  in  nitrous  oxide  are  the 
same  as  in  pure  oxygen,  showing  that  the  compound  N20  is 
readily  broken  down,  yielding  free  oxygen. 

When  nitrous  oxide  is  inhaled  into  the  lungs,  it  produces  a 
remarkable  physiological  effect,  generally  throwing  the  subject 
into  a  hysterical  condition.  It  is,  therefore,  known  as  laughing 
gas.  When  consumed  in  larger  quantity  it  produces  anesthesia, 
and  is  consequently  used  in  minor  surgical  operations. 

Nitrous  oxide  is  a  colorless  gas,  with  a  sweetish  taste,  and 
dissolves  readily  in  cold  water.  It  should,  therefore,  be  col- 
lected in  cylinders  over  hot  water. 

When  most  chemical  reactions  take  place,  heat  is  evolved,  — 
most  reactions  are  exothermic.  In  this  case  the  opposite  is 
true.  When  nitrogen  and  oxygen  combine  to  form  nitrous 
oxide,  heat  is  absorbed.  Such  chemical  reactions,  which  take 
place  with  absorption  of  heat,  are  known  as  endothermic  re- 
actions. 

Nitric  Oxide,  NO.  —  Nitric  oxide  can  be  formed  by  the  reduc- 
tion of  the  higher  oxides  of  nitrogen  or  by  the  reduction  of 
nitrous  or  nitric  acid.  It  is  prepared  most  conveniently  by 
the  action  of  nitric  acid  on  metallic  copper.  The  equations 
expressing  the  reaction  are  :  — 


3  CuO  +  6  HN03  =  3  Cu  (N03)  2  +  3  H20. 

As  quickly  as  the  colorless  gas,  nitrous  oxide,  is  brought  in 
contact  with  free  oxygen,  the  two  combine  at  ordinary  temper- 
ature:— 


=  2N0 


The  gas  N02  has  a  yellowish  brown  color.  These  are  the 
fumes  which  always  appear  when  nitric  acid  acts  on  metallic 
copper.  Nitric  oxide  produces  a  dark  violet  color  when 


82  ELEMENTS  OF  INORGANIC  CHEMISTRY 

brought  in  contact  with  a  solution  of  a  ferrous  salt.  This 
reaction  is  used  to  detect  nitric  oxide. 

Nitrogen  Sesquioxide  or  Nitrogen  Trioxide,  N203.  —  Nitrogen 
sesquioxide  is  obtained  by  the  altion  of  arsenic  trioxide,  As203, 
upon  nitric  acid.  Also  by  the  action  of  a  strong  acid,  like  sul- 
phuric, upon  nitrites  :  — 

2  NaN02  +  H2S04  =  Na,S04  +  H,0  +  N203- 

It  is,  therefore,  sometimes  called  nitrous  anhydride,  since  it  is 
nitrous  acid  minus  water,  and,  by  the  addition  of  an  alkali,  it 
forms  nitrites.  Nitrogen  sesquioxide  is  stable  only  at  low 
temperatures.  Above  —  20°,  or  —  15°,  it  begins  to  decompose 
into  nitrogen  dioxide  and  nitric  oxide  :  — 


At  very  low  temperatures  it  passes  over  into  a  deep  blue  liquid. 
Nitrogen  Dioxide  or  Nitrogen  Peroxide,  N02.  —  Nitrogen 
dioxide  is  conveniently  formed  by  the  action  of  oxj^gen  on 
nitric  oxide  :  — 


Nitrogen  peroxide  is  a  strong  oxidizing  agent.     It  is  also  very 
poisonous. 

Nitrogen  Pentoxide,  N205.  —  Nitrogen  pentoxide  is  formed  by 
the  action  of  strong  dehydrating  agents,  like  phosphorus  pen- 
toxide, upon  nitric  acid  :  — 

2  HN03  +  P205  =  (P205.H20)  +N205. 

Nitrogen  pentoxide  readily  combines  with  water,  forming  nitric 
acid  :  — 


It  is,  therefore,  the  anhydride  of  nitric  acid.  It  is  a  powerful 
oxidizing  agent,  as  would  be  expected  from  the  large  amount 
of  oxygen  which  it  contains. 

Nitrous  Acid,  HN02.  —  The  salts  of  nitrous  acid,  the  nitrites, 
are  obtained  by  removing  oxygen  from  the  nitrates. 


NITROGEN  83 

If  a  mild  reducing  agent,  such  as  metallic  lead,  is  fused  with 
a  nitrate,  the  reduction  to  nitrite  takes  place  far  more  easily 
and  completely  :  — 

KN03  +Pb  =*PbO  +  KN02. 

When  a  nitrite  is  treated  with  a  strong  acid,  such  as  sul- 
phuric, nitrous  acid  is  set  free  :  — 

2  K£K)2  +  H2S04  =  K2S04  +  2  HK02. 


Nitrous  acid  can  exist  only  in  solution.  When  an  attempt 
is  made  to  remove  the  water,  the  nitrous  acid  loses  water  and 
passes  into  the  anhydride,  N203  :  —  v 


Nitrous  acid  is  an  excellent  reducing  agent,  since  it  readily 
combines  with  oxygen,  forming  nitric  acid.  When  brought  in 
contact  with  a  substance  rich  in  oxygen,  like  potassium  per- 
manganate, it  takes  oxygen  away  from  the  compound,  convert- 
ing it  into  colorless  substances.  The  destruction  of  the 
beautiful  purple  color  takes  place  very  rapidly. 
X  Nitric  Acid,  HN03.  —  This  is  not  only  the  most  important 
acid  of  nitrogen,  but  one  of  the  strongest  and  most  important 
of  all  known  acids.  It  was  early  prepared  from  nitre,  which 
is  potassium  nitrate,  whence  its  name. 

H2S04  +  2  KN03  =  K2S04  +  2  HN03. 

Potassium  nitrate,  known  as  saltpetre,  is  formed  where 
organic  matter  is  decomposing  in  the  presence  of  potassium 
salts.  It  occurs  in  the  form  of  a  solid  only  in  arid  regions, 
since,  on  account  of  its  great  solubility,  it  would  pass  into  solu- 
tion if  it  came  in  contact  with  any  appreciable  amount  of 
water.  Sodium  nitrate  occurs  in  abundance  in  the  arid  regions 
of  Chili,  and  is  known  as  Chili  saltpetre.  When  sodium 
nitrate  is  treated  with  sulphuric  acid,  the  first  reaction  that 
takes  place  is  :  — 

4  =  NaHS04  +  HN03. 


84  ELEMENTS  OF  INORGANIC   CHEMISTRY 

If  the  temperature  is  raised  sufficiently,  the  acid  sulphate  acts 
on  more  of  the  nitrate,  decomposing  it  in  the  sense  of  the 
following  equation :  — 

NaHS04  +  NaN03  =  Na2S04  +  HN03. 

The  decomposition  of  nitrates  by  sulphuric  acid  does  not 
mean  that  sulphuric  acid  is  stronger  than  nitric  acid.  The 
nitric  acid  being  volatile  is  set  free,  although  it  is  much 
stronger  than  sulphuric  acid. 

Chemical  Properties  of  Nitric  Acid.  —  The  most  characteristic 
chemical  property  of  nitric  acid  is  its  strong  oxidizing  power. 
When  brought  in  contact  with  substances  which  can  take  up 
oxygen,  nitric  acid  readily  gives  up  its  oxygen  and  passes  over 
into  lower  oxides  of  nitrogen. 

The  oxidizing  action  of  nitric  acid  and  its  salts  can  be  repre- 
sented as  follows :  — 

4  HN03  =  2  H20  +  4  N02  +  02 ; 
4  HN03  =  2  H20  +  4  NO  +  3  O2; 
=  2H20+2N2  +  502; 


the  reaction  that  takes  place  in  any  given  case  depending  upon 
the  nature  of  the  substance  present  and  upon  the  conditions. 

When  a  metal  is  treated  with  nitric  acid,  the  hydrogen  ion 
of  the  acid  gives  up  its  charge  to  the  metal,  converting  the 
latter  into  an'  ion,  while  the  hydrogen  becomes  an  atom.  The 
hydrogen,  in  the  case  of  nitric  acid,  however,  does  not  escape, 
but  acts  on  more  nitric  acid,  reducing  it  to  lower  oxides  of 
nitrogen,  or  to  nitrogen  itself,  or  even  to  ammonia,  depending 
upon  conditions. 

Nitric  acid  is  a  liquid,  boiling  with  partial  decomposition  at 
86°. 

Nitric  acid  and  water  are  miscible  in  all  proportions. 

Detection  of  Nitric  Acid. — Nitric  acid  is  readily  detected 
by  the  dark  purple  color  produced  when  it  is  mixed  with 
a  concentrated  solution  of  ferrous  sulphate,  both  solutions 
being  warm.  The  test  for  nitric  acid  is  made  as  follows :  The 


NITROGEN  85 

nitric  acid  or  the  nitrate  is  treated  with  a  little  concentrated 
sulphuric  acid,  and  warmed  until  the  containing  vessel  feels 
warm  to  the  hand.  Another  test-tube  is  filled  about  one-third 
full  of  crystals  of  ferrous  sulphate,  and  dissolved  in  just  as 
little  water  as  possible,  the  solution  being  heated  until  it  feels 
warm  to  the  hand,  but  not  heated  to  boiling.  The  solution 
containing  the  nitric  acid  is  now  added  drop  by  drop  to  the 
solution  of  ferrous  sulphate,  when  the  dark  color  will  make  its 
appearance  in  the  form  of  a  ring  where  the  two  liquids  come 
in  contact. 

Dissociation  of  Nitric  Acid  and  Nitrates.  —  Nitric  acid  dis- 
sociates in  the  sense  of  the  following  equation :  — 

HN03  =  H,  N~03. 

It  is,  therefore,  a  monobasic  acid,  and  can  yield  only  one  series 
of  salts.  These  are  of  the  general  type,  MN03,  and  dissociate 
thus : — 

MN03  =  M,  N03. 

Fuming  Nitric  Acid.  —  Fuming  nitric  acid  is  formed  in  the 
preparation  of  nitric  acid  from  sodium  nitrate  and  sulphuric 
acid,  if  the  temperature  is  sufficiently  high.  It  is  apparently 
a  solution  of  nitrogen  dioxide  in  nitric  acid. 

It  is  a  much  more  energetic  oxidizing  agent  than  ordinary 
concentrated  nitric  acid.  When  warmed  in  its  fumes  many 
organic  substances  take  fire  and  burn. 

Aqua  Regia. — Certain  metals,  like  gold  and  platinum,  do 
not  dissolve  in  nitric  acid,  but  when  treated  with  a  mixture  of 
nitric  and  hydrochloric  acids  they  dissolve  readily.  The  mix- 
ture which  is  most  efficient  consists  of  one  part  of  nitric  acid 
and  three  parts  of  hydrochloric  acid.  This  is  known  as  aqua 
regia.  The  nitric  acid  in  the  mixture  oxidizes  the  hydrochloric 
acid  and  liberates  chlorine.  There  is  probably  also  formed 
one  or  more  compounds  containing  nitrogen,  oxygen,  and 
chlorine. 

The  action  of  these  various  substances   is   to  convert   the 


86 


ELEMENTS  OF  INORGANIC  CHEMISTRY 


metals  into  chlorides,  even  platinum  being  transformed  into 
platinic  chloride  by  aqua  regia.  The  name  was  derived  from 
the  fact  that  this  mixture  can  dissolve  gold. 


EXPERIMENTS  WITH  NITROGEN  AND  ITS  COMPOUNDS 
WITH  HYDROGEN  AND  OXYGEN 

Experiment  38.  Nitrogen  obtained  from  the  Air.  —  (Bell-jar 
closed  by  a  cork;  pneumatic  trough  tilled  with  water;  small  por- 
celain dish ;  a  piece  of  phosphorus  the  size  of  a  pea.) 

Arrange  the  apparatus 
as  shown  in  Figure  24.  A 
small  piece  of  phosphorus, 
not  larger  than  a  pea,  is 
dried  by  pressing  it  be- 
tween folds  of  filter-paper. 
Phosphorus  must  never  be 
touched  with  the  hand,  or 
painful  wounds  may  result 
from  the  ignition  of  the 
phosphorus.  The  phos- 
phorus is  placed  in  the 
porcelain  dish,  which  is 
floated  upon  water  as  in 
the  figure.  The  phos- 
phorus is  ignited,  and  the  bell-jar  placed  over  it  on  the  shelf 
in  the  trough,  so  that  it  dips  well  beneath  the  surface  of  the 
water.  As  the  phosphorus  burns  the  air  in  the  bell-jar  becomes 
heated,  and  some  may  escape  in  the  form  of  bubbles  from  the 
bottom  of  the  jar.  After  the  phosphorus  has  ceased  to  burn, 
the  water  will  rise  in  the  bell-jar. 

Examine  the  gas  which  remains  in  the  jar  to  see  whether  it 
is  nitrogen.  After  bringing  the  water  inside  the  jar  to  the 
same  level  as  that  outside,  by  raising  or  lowering  the  jar  in  the 
trough,  remove  the  cork  and  introduce  quickly  a  burning 
taper,  or  a  burning  piece  of  sulphur,  and  observe  what  happens. 
Nitrogen  will  not  support  combustion.  Do  these  results  indi- 
cate the  presence  of  nitrogen?  (See  page  76.) 

Experiment  39.  Preparation  of  Nitrogen  from  Ammonium 
Nitrite.  —  (Flask  holding  250  to  300  cc. ;  thistle-tube ;  two-hole 
stopper ;  exit-tube ;  ammonium  chloride ;  sodium  nitrite.) 


FIG.  24. 


NITROGEN 


87 


Introduce  into  the  flask  (Fig.  25)  15 
grams  of  ammonium  chloride  and  20 
grams  of  sodium  nitrite.  Pour  in  50  cc. 
of  water.  Heat  the  flask  slowly.  Nitro- 
gen. will  be  evolved  before  there  is  any 
boiling  in  the  flask.  When  ammonium 
chloride  is  heated  in  such  a  solution 
with  sodium  nitrite,  ammonium  nitrite 
is  formed  in  sense  of  the  following 
equation  :  — 


NH4C1 


=  NaCl  +  NH4N02. 


The  ammonium  nitrite  then  decom- 
poses in  the  manner  already  described. 

Test  the  gas  as  in  the  last  experi- 
ment. Does  it  support  combustion  ? 
(See  page  76.) 

Experiment  40.  Preparation  of  Am- 
monia from  Ammonium  Chloride  and 
Sodium  Hydroxide.  —  (Test-tube  ;  am- 
monium chloride;  sodium  hydroxide.) 

Introduce  into  a  test-tube  enough  am- 


FIG.  25. 


monium  chloride  to  make  a  layer  about  1  cm.  deep.    Add  a  few 
drops  of  a  concentrated  solution  of  sodium  hydroxide.     Notice 

the  odor  of  the  gas  that 
escapes.  Warm  the  tube 
gently  and  notice  that  the 
odor  becomes  more  intense. 
Moisten  a  piece  of  red  lit- 
mus paper  and  observe  the 
change  in  color.  This  is  a 
characteristic  reaction  of 
alkalies  in  general. 
(See  page  77.) 

Experiment  41. 
Preparation  of  Am- 
monia from  Ammo- 
nium Chloride  and 
Slaked  Lime.  — 
(Flask  holding  500 
cc. ;  stopper;  deliv- 
ery-tube ;  funnel  j 
FIG.  26.  beaker.) 


88 


ELEMENTS   OF  INORGANIC   CHEMISTRY 


Arrange  the  apparatus  as  shown  in  Figure  26.  Into  the 
flask  A  introduce  150  grams  of  dry  slaked  lime  mixed  with 
100  grams  of  ammonium  chloride.  Place  the  flask  upon  a 
sand-bath  or  an  asbestos  board  and  heat  gently.  Ammonia  gas 
will  escape.  Collect  some  of  the  gas  in  water  by  attaching  a 
funnel  F  to  the  delivery-tube  as  in  the  figure  and  as  in  Figure 
23.  The  funnel  must  dip  only  a  short  distance  beneath  the 
surface  of  the  water.  Ammonia 
is  very  soluble  in  water,  and  if 
this  precaution  is  not  taken  the 
water  will  absorb  the  ammonia 
very  rapidly  and  rush  back  into 
the  flask  A.  This  solution,  known 
as  ammonia  water,  should  be  pre- 
served for  future  experiments. 

Collect  some  of  the  ammonia 
in  cylinders  by  attaching  the  tube 
T  (Fig.  27)  to  the  delivery-tube 
from  the  flask  A  (Fig.  26).  The 
tube  T  must  extend  to  the  top  of 
the  inverted  cylinder,  since  am- 
monia is  lighter  than  air.  The 
ammonia  then  displaces  the  air  and 
the  cylinder  can  be  readily  filled 
with  the  gas.  (See  page  77.) 

Experiment  42.     Ammonia  dis- 
solved in  Water  is  a  Base. 
—  (Aqueous    solution    of 
ammonia;     litmus;     phe- 
nolphthalein.) 

To  a  portion  of 

the    solution    of   ( 

ammonia  in  water 
obtained   in   EX- 


FIG.  27. 


pertinent  41,  add  a  few  drops  of  litmus.  Litmus  in  the  pres- 
ence of  a  neutral  substance  is  purple.  In  the  presence  of 
aqueous  ammonia  it  is  deep  blue  in  color.  This  reaction  is 
characteristic  of  alkalies  in  general. 

To  another  portion  of  the  aqueous  solution  of  ammonia  add 
a  few  drops  of  the  alcoholic  solution  of  phenolphthale'in.  The 
solution  becomes  deep  red  in  color.  This  reaction  is  also  char- 
acteristic of  alkalies.  This  color  may  disappear  in  the  case  of 
ammonia  after  a  time.  (See  page  78.) 


NITROGEN  89 

Experiment  43.     Neutralization  of  Ammonia  by  Acids. — 

(Three  beakers;  aqueous  solution  of  ammonia;  dilute  hydro- 
chloric acid ;  dilute  sulphuric  acid ;  dilute  nitric  acid.) 

Introduce  50  cc.  of  distilled  water  into  each  beaker  and  add 
25  cc.  of  the  solution  of  ammonia  to  each.  Also  a  few 
drops  of  litmus. 

Into  one  beaker  drop  from  a  pipette  dilute  hydrochloric  acid, 
stirring  the  solution  vigorously  by  means  of  a  stirring  rod  each 
time  a  drop  of  the  acid  falls  into  the  solution  of  ammonia. 
After  a  time  the  solution  will  turn  red,  showing  that  all  the 
ammonia  has  been  neutralized. 

Add  sulphuric  acid  in  the  same  manner  to  the  second  beaker, 
and  nitric  acid  to  the  third  until  the  litmus  turns  red  in  each 
case.  Preserve  the  solutions.  (See  page  80.) 

Experiment  44.  A  Salt  is  formed  when  Ammonia  is  Neutral- 
ized by  an  Acid.  —  (Three  small  evaporating  dishes ;  the  three 
solutions  obtained  in  the  last  experiment.) 

Pour  the  three  solutions  obtained  in  the  last  experiment  into 
the  three  evaporating  dishes  —  one  in  each  dish.  Evaporate 
the  solutions  to  dryness.  A  white  salt  will  remain  in  each 
dish  —  ammonium  chloride  in  one  dish,  ammonium  sulphate  in 
another  dish,  and  ammonium  nitrate  in  the  third. 

Introduce  some  of  the  ammonium  chloride  into  a  test-tube, 
some  of  the  ammonium  sulphate  into  a  second  test-tube,  and 
some  of  the  ammonium  nitrate  into  a  third  test-tube.  Add  a 
few  drops  of  caustic  soda  to  each  tube  and  warm  gently. 
Notice  the  odor  of  ammonia  escaping  from  each  tube.  Test 
for  ammonia  with  red  litmus  paper.  This  reaction  is  character- 
istic of  ammonium  salts.  (See  page  80.) 

Experiment  45.  Gaseous  Ammonia  reacts  with  Gaseous 
Hydrochloric  Acid.  —  (Cylinder  filled  with  ammonia;  two  glass 
plates ;  cylinder  containing  a  few  drops  of  concentrated  hydro- 
chloric acid.) 

Introduce  a  few  drops  of  concentrated  hydrochloric  acid 
into  a  glass  cylinder  and  cover  the  top  of  the  cylinder  with 
a  glass  plate.  Place  above  this  a  cylinder  filled  with  am- 
monia gas  also  covered  with  a  glass  plate.  *  The  apparatus 
should  be  arranged  as  in  Figure  13.  Eemove  the  plates, 
bringing  the  two  cylinders  together.  The  gases  combine  at  once, 
forming  ammonium  chloride,  which  is  deposited  upon  the  walls 
as  a  white  powder.  Collect  a  little  of  the  powder,  place  it  in  a 
test-tube,  and  add  a  few  drops  of  sodium  hydroxide.  Notice  the 
odor  of  ammonia.  This  characteristic  reaction  of  ammonium 


90 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


salts  shows  that  a  salt  of  ammonium  was  formed  in  the  above 
reaction.     (See  page  78.) 

Experiment  46.  Ammonia  Lighter  than  Air.  —  (Large  bal- 
ance (Fig.  2) ;  two  large  beakers ;  large  bottle  filled  with 
ammonia.) 

Eemove  the  glass  cylinders  from  the  large  balance  (Fig.  2), 
and  suspend  a  large  beaker  from  each  arm  of  the  balance. 
When  equilibrium  has  been  established  and  the  pointer  made 
to  stand  at  zero  by  adding  weights  to  one  or  the  other  arm,  the 
bottle  of  ammonia  is  opened  beneath  one  of  the  beakers. 
This  arm  of  the  balance  will  be  seen  to  rise  at  once,  showing 
that  ammonia  is  lighter  than  air.  (See  page  78.) 

Experiment  47.  Absorption  of  Ammonia  by  Water.  —  (Large 
glass  tube  about  30  cm.  in  length  closed  at  one  end ;  iron  stand 
and  clamp ;  mercury  trough.) 

Fill  the  glass  tube  (Fig.  28)  with  mercury  and  insert  it  into 
the  mercury  trough,  which  also  contains  considerable  mercury. 

Fill  the  tube  with  ammonia  by  lead- 
ing the  discharge  tube  from  the  am- 
monia-generating apparatus  beneath 
the  lower  end  of  the  glass  tube.  As  the 
ammonia  rises  in  the  glass  tube,  the 
mercury  will  be  displaced  and  will 
flow  out  into  the  mercury  trough. 

When  the  glass  tube  has  become 
filled  with  ammonia  gas,  add  a  few 
drops  of  water  to  the  tube  by  means 
of  a  bent  glass  tube  filled  with 
water  and  placed  beneath  the  lower 
end  of  the  tube  containing  ammonia. 
The  water  will  rise  to  the  top  of  the 
mercury  in  the  tube,  absorb  the  am- 
monia, and  the  mercury  will  rise  in 
the  tube  as  the  ammonia  is  absorbed. 
(See  page  78.) 

Experiment   48.      Absorption  of 
Ammonia  by  Water. — (Two  large 
balloon   flasks;    glass  tubing;    one 
one-hole  rubber  stopper;   one  two- 
FlG  2g  hole  rubber  stopper ;  iron  stand  and 

clamp.) 

The  apparatus  is  to  be 'arranged  as  in  Figure  29.  The  balloon 
flask  B,  holding  about  2  litres,  is  cleaned  and  dried  and  filled 


NITROGEN 


91 


with  ammonia  by  inverting  the  flask  and  passing  the  escape- 
tube  from  the  ammonia-generating  flask  up  to  the  top  of  the 
inverted  flask.  Ammonia  being  lighter  than  air  drives  the  air 
downwards  out  of  the  flask  and  completely  fills  it.  The  lower 
flask  At  holding  3  litres,  is  filled 
with  distilled  water  containing  a 
little  litmus,  to  which  just  enough 
acid  is  added  to  color  the  litmus 
red. 

The  apparatus  is  now  arranged 
as  shown  in  the  figure.  The  breath 
is  then  blown  into  the  tube  G. 
This  forces  a  little  water  from  the 
flask  A,  through  the  tube  D,  into 
the  flask  B.  The  water  which 
enters  B  absorbs  the  ammonia, 
producing  a  partial  vacuum  in  B. 
Water  immediately  rushes  up  from 
A  into  B,  and  a  fountain  results. 
The  ammonia  in  A  first  neutral- 
izes the  acid  in  the  water  and  then 
renders  the  water  alkaline,  chang- 
ing the  color  of  the  litmus  from 
red  to  blue. 

If  the  precaution  was  taken  to 
completely  fill  the  flask  B  with 
ammonia,  i.e.  to  pass  in  the  am- 
monia until  all  the  air  had  been 
displaced,  water  will  rise  in  B  until 
this  flask  is  practically  full.  FlG  29 

The  iron  stand  and  clamp  must 

be  so  firm  as  to  hold  the  flask  B  when  full  of  water.     (See 
page  78.) 

Experiment  49.  Absorption  of  Ammonia  by  Charcoal. — 
(Glass  tube  filled  with  ammonia ;  mercury  trough ;  a  piece  of 
charcoal  which  has  been  freshly  heated.) 

Fill  a  glass  tube  with  ammonia,  as  in  Figure  28.  Introduce 
a  piece  of  freshly  heated  charcoal  by  pushing  it  under  the 
mercury  beneath  the  mouth  of  the  tube,  and  allowing  it  to 
rise  in  the  tube.  The  charcoal  will  absorb  the  ammonia  and 
the  mercury  will  rise  in  the  tube. 

This  illustrates  the  way  in  which  charcoal  absorbs  ammonia 
and  other  gases  from  wa.te£  anci  thus  purifies  it. 


92  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Experiment  50.  Preparation  of  Nitrous  Oxide.  —  (Ketort 
holding  about  100  cc. ;  pneumatic  trough ;  glass  cylinders  with 
glass  plates  as  covers ;  ammonium  nitrate.) 

Introduce  15  grams  of  ammonium  nitrate,  which  had  been 
previously  dried  at  110°  in  an  air-bath  to  remove  moisture, 
into  the  retort  arranged  as  in  Figure  3.  Heat  the  retort  very 
gently  ;  indeed,  just  rapidly  enough  to  obtain  a  steady  current 
of  gas.  Rapid  heating  may  produce  frothing  in  the  retort,  or 
even  explosion. 

It  is  better  to  have  the  water  in  the  trough  somewhat  warm, 
since  nitrous  oxide  is  less  soluble  in  warm  water  than  in  cold. 

Fill  several  cylinders  with  the  gas  for  use  in  the  next 
experiment.  (See  page  80.) 

Experiment  51.  Nitrous  Oxide  supports  Combustion.  —  (Four 
or  five  cylinders  filled  with  nitrous  oxide ;  deflagrating-spoon ; 
a  splinter;  sulphur;  phosphorus;  steel  watch-spring.) 

Into  one  of  the  cylinders  introduce  a  splinter  which  has 
been  burned  at  the  end  to  a  coal,  but  from  which  the  flame 
has  been  extinguished.  It  will  be  rekindled  into  flame. 

Into  another  cylinder  introduce  highly  heated,  burning  sul- 
phur in  a  deflagrating-spoon,  and  into  still  another  cylinder 
introduce  burning  phosphorus.  They  will  burn  with  intense 
light  and  with  the  evolution  of  a  large  amount  of  heat,  as  in 
oxygen. 

Cover  the  bottom  of  the  fourth  cylinder,  filled  with  nitrous 
oxide,  with  a  sheet  of  asbestos  and  a  layer  of  water,  and  plunge 
into  the  cylinder  a  steel  watch-spring  tipped  with  sulphur  and 
ignited.  The  iron  will  burn  very  vigorously,  sending  off  brill- 
iant scintillations.  The  asbestos  and  water  were  introduced 
into  the  cylinder  so  that  the  molten  metal  and  oxide  would  not 
fall  on  to  the  bottom  of  the  cylinder  and  break  it. 

In  all  of  the  above  experiments  the  products  of  combustion 
are  the  same  as  in  oxygen,  and,  therefore,  they  suggest  com- 
bustion in  oxygen  in  a  most  striking  manner.  (See  page  81.) 

Experiment  52.  Preparation  of  Nitric  Oxide.  —  (A  balloon 
flask,  or  Florence  flask,  holding  500  cc. ;  a  thistle-tube  ;  a  de- 
livery-tube bent  at  right  angles;  a  two-hole  stopper;  copper 
turnings ;  concentrated  nitric  acid.) 

Into  the  balloon  flask  (Fig.  30)  introduce  some  copper  turn- 
ings and  pour  slowly  on  to  the  turnings  concentrated  nitric  acid. 
When  a  steady  stream  of  the  gas  is  evolved  cease  to  add  the 
nitric  acid,  since  an  excess  of  acid  frequently  produces  frothing 
in  the  flask. 


NITROGEN 


93 


The  gas  which  first  escapes  is  reddish  brown  in  color,  due  to 
the  combination  of  nitric  oxide  with  the  oxygen  of  the  air  in 
the  flask,  forming  nitric  oxide. 

After  the  gas  in  the  flask  has  become 
practically  colorless,  fill  several  cylin- 
ders in  a  pneumatic  trough  over  water. 
The  gas  in  the  cylinders  will  be  seen  to 
have  very  little  color.  (See  page  81.) 

Experiment  53.  Combination  of  Nitric 
Oxide  with  Oxygen. —  (Glass  cylinder 
filled  with  nitric  oxide,  covered  with  a 
glass  plate;  glass  cylinder  filled  with  air.) 

Arrange  the  apparatus  as  in  Figure 
13 ;  remove  the  glass  plates  between 
the  two  cylinders  and  allow  the  nitric 
oxide  to  come  in  contact  with  the  air. 
Dense  red  fumes  will  appear  imme- 
diately on  bringing  the  two  gases  in 
contact.  Observe  carefully  the  move- 
ments of  the  red  vapor  in  the  cylinder. 
(See  page  81.) 

Experiment  54.  Nitric  Oxide  does 
not  support  Combustion.  —  (Three  or 
four  cylinders  filled  with  nitric  oxide ; 
a  splinter;  a  piece  of  charcoal;  a  piece  FlG  30 

of  sulphur.) 

Into  a  cylinder  filled  with  nitric  oxide  introduce  a  burning 
splinter ;  observe  what  occurs. 

Into  another  cylinder  filled  with  nitric  oxide  introduce  a 
piece  of  sulphur  ignited  in  a  deflagrating-spoon.  Does  it  con- 
tinue to  burn  ? 

Into  a  third  cylinder  filled  with  the  gas  introduce  a  piece  of 
carbon  heated  to  redness.  Does  the  carbon  continue  to  burn  ? 

Compare  these  results  with  those  obtained  with  nitrous 
oxide. 

Experiment  55.  Preparation  of  the  Anhydride  of  Nitrous 
Acid;  Nitrogen  Sesquioxide.  —  (Balloon  or  Florence  flask 
holding  200  to  300  cc. ;  thistle-tube;  two-hole  stopper;  delivery- 
tube  bent  at  right  angles;  small  cylinder;  large  bottle  or 
cylinder;  a  U-tube  surrounded  by  ice  and  salt;  arsenic  tri- 
oxide ;  nitric  acid  of  the  specific  gravity  1.35.) 

Thirty  grams  of  arsenic  trioxide  are  introduced  into  the 
flask  A  (Fig.  31).  This  is  covered  with  nitric  acid  of  the 


94 


ELEMENTS   OF  INORGANIC   CHEMISTRY 


specific  gravity  1.35,  and  gently  warmed.  The  heavy  red  gas, 
nitrogen  trioxide,  can  be  collected  in  a  bottle  or  cylinder 
by  leading  the  exit-tube  to  the  bottom  of  the  vessel  placed 
upright. 

If  the  gas  is  passed  through  the  U-tube  (Fig.  31)  surrounded 
by  ice  and  salt,  it  condenses,  forming  a  blue  liquid.     The  gas  is 
first  passed  through  the  bottle  B  in  order  that  any 
steam  may  condense.     (See  page  82.) 

Experiment  56.  Preparation  of  Nitrogen  Peroxide. 
—  (Balloon  or  Florence  flash  holding  250  cc. ;  two- 
hole  rubber  stopper;  thistle-tube;  discharge-tube; 
U-tube ;  tin ;  concentrated  nitric  acid  ;  ice ;  salt.) 


u 


FIG.  31. 


Introduce  into  the  flask  some  granulated  tin,  and  cover  this 
with  concentrated  nitric  acid.  The  gas  may  be  passed  through 
an  empty  flask  to  allow  the  steam  to  condense,  and  then 
through  the  U-tube  surrounded  by  ice  and  salt  as  in  Figure  31. 
The  gas  is  liquefied  in  the  U-tube,  forming  a  reddish  liquid. 

Compare  this  with  the  liquid  obtained  in  Experiment  55. 
(See  page  82.) 

Experiment  57.  Preparation  of  Nitric  Acid.  —  (Tubulated 
retort  holding  from  300  to  500  cc. ;  receiver ;  large  funnel ; 
sodium  nitrate  ;  concentrated  sulphuric  acid.) 

Arrange  the  apparatus  as  in  Figure  32.  Introduce  into  the 
retort  50  grams  of  sodium  nitrate,  and  add  25  grams  of  concen- 
trated sulphuric  acid.  Warm  gently  the  contents  of  the. 


NITROGEN 


95 


retort.  Nitric  acid  will  condense  in  the  receiver,  which  is 
kept  cool  by  a  current  of  water  which  flows  over  the  outside 
of  the  receiver.  Nitric  acid  prepared  in  this  way  is  colored 
somewhat  yellow,  due  to  the  presence  of  oxides  of  nitrogen 
dissolved  in  it.  With  nitric  acid  perform  the  following  experi- 
ments. (See  page  83.) 


FIG.  32. 

Experiment  58.  Acid  Properties  of  Nitric  Acid.  —  (Nitric 
acid ;  sodium  hydroxide ;  litmus.) 

Dip  a  piece  of  blue  litmus  paper  into  a  solution  of  a  few 
drops  of  nitric  acid  in  a  large  volume  of  water.  Observe  how 
quickly  the  blue  litmus  turns  deep  red  in  color  —  a  character- 
istic reaction  of  acids. 

1  Introduce  10  cc.  of  concentrated  nitric  acid  into  a  beaker 
containing  50  cc.  of  distilled  water;  note  the  temperature;  add 
a  few  drops  of  litmus.  Add  cautiously  from  a  pipette  potas- 
sium hydroxide,  stirring  the  solution  vigorously.  As  soon  as 
the  litmus  changes  from  red  to  blue,  cease  to  add  the  alkali. 
Note  again  the  temperature  of  the  solution.  Has  there  been 
any  change  ?  Why  ? 

Evaporate  the  solution  to  dryness ;  a  white  solid  is  obtained 
which  is  potassium  nitrate.  Save  for  future  use.  (See  page  84.) 


96  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Experiment  59.  Fuming  Nitric  Acid.  —  (Tubulated  retort 
holding  300  cc.  to  500  cc. ;  receiver ;  concentrated  nitric  acid, 
concentrated  sulphuric  acid.) 

This  experiment  is  to  be  performed  only  by  the  instructor. 
Arrange  the  apparatus  as  in  Figure  32.  Pour  cautiously  150 
grams  of  concentrated  sulphuric  acid  into  75  grams  of  con- 
centrated nitric  acid.  Introduce  the  mixture  into  the  retort, 
and  distil  sloivly  under  the  hood.  The  nitric  acid  which  con- 
denses and  is  caught  in  the  receiver  is  known  as  fuming  nitric 
acid.  The  sulphuric  acid  has  taken  up  a  part  of  the  water 
present,  and  the  nitric  acid  which  distils  over  is  much  stronger 
than  ordinary  concentrated  nitric  acid.  (See  page  85.) 

Experiment  60.  Oxidizing  Power  of  Nitric  Acid.  —  (Action 
on  sugar.)  (Nitric  acid  ;  Erlenmeyer  flask  100  cc. ;  cane  sugar.) 

Introduce  5  grams  of  cane  sugar  into  the  Erlenmeyer  flask, 
and  add  20  cc.  of  concentrated  nitric  acid.  Little  or  no  reac- 
tion takes  place.  Warm  the  mixture.  A  violent  reaction  will 
take  place,  resulting  in  the  evolution  of  large  volumes  of  dense 
red  fumes.  The  sugar  is  oxidized  by  the  hot  nitric  acid. 
Evaporate  the  solution  until  only  a  small  volume  remains. 
Set  this  aside  in  a  beaker  or  crystallizing  dish,  and  beautiful 
white  crystals  will  separate.  These  are  crystals  of  oxalic  acid. 
(See  page  84.) 

Experiment  61.  Oxidizing  Power  of  Nitric  Acid.  —  (Action 
on  wood  and  charcoal.)  (An  Erlenmeyer  flask  100  cc. ;  beaker 
holding  200  cc. ;  fuming  nitric  acid ;  splinter ;  stick  of  charcoal.) 

Introduce  25  cc.  of  the  acid  into  the  Erlenmeyer  flask  and 
heat.  Plunge  a  glowing  splinter  into  the  liquid.  It  will  burn 
vigorously. 

Pour  50  cc.  of  fuming  nitric  acid  into  the  beaker.  Heat 
one  end  of  the  stick  of  charcoal  to  redness,  holding  the  stick 
with  a  pair  of  forceps.  Quickly  plunge  the  hot  end  into  the 
nitric  acid.  The  charcoal  will  continue  to  burn  under  the 
liquid.  Where  does  the  oxygen  which  combines  with  the  car- 
bon come  from  ? 

This  experiment  must  be  performed  under  the  hood  on  ac- 
count of  the  escape  of  noxious  vapors.  (See  page  84.) 

Experiment  62.  Oxidizing  Power  of  Nitric  Acid.  —  (Action 
on  tin.)  (Erlenmeyer  flask;  concentrated  nitric  acid;  some 
granules  of  tin.) 

Into  a  small  Erlenmeyer  flask  introduce  a  few  granules  of 
tin;  cover  these  with  concentrated  nitric  acid  and  warm 
cautiously.  Eed  fumes  will  escape,  and  the  tin  will  disappear. 


NITROGEN  97 

It  will,  however,  not  pass  into  solution,  but  a  white  powder 
will  be  formed.  This  is  an  oxide  of  tin. 

"From  what  source  did  the  tin  obtain  the  oxygen?  (See 
page  84.) 

Experiment  63.  Nitric  Acid  dissolves  Certain  Metals.  —  (Er- 
lenmeyer  flask  100  cc. ;  granules  of  zinc ;  nitric  acid.) 

Into  the  Erlenmeyer  flask  introduce  some  pieces  of  zinc  and 
cover  the  zinc  with  nitric  acid.  The  zinc  rapidly  passes  into 
solution.  What  escapes  ?  Evaporate  the  solution  to  dryness. 
White  zinc  nitrate  separates.  (See  page  84.) 

Experiment  64.  Action  of  Nitric  Acid  on  Hydrochloric  Acid 
(Aqua  Eegia).  —  (Large  test-tube ;  concentrated  nitric  acid ; 
concentrated  hydrochloric  acid.) 

Mix  10  cc.  of  concentrated  nitric  acid  and  30  cc.  of  concen- 
trated hydrochloric  acid,  and  fill  the  test-tube  one-third  full  of 
the  mixture.  Warm  gently  over  a  bunsen  burner.  A  gas  will 
escape  whose  odor  suggests  that  of  chlorine.  It  contains  in 
addition  to  chlorine  certain  oxides  of  nitrogen  and  chlorine. 

The  nitric  acid  oxidizes  the  hydrochloric  acid. 

This  mixture  readily  dissolves  gold  and  even  platinum. 
(See  page  85.) 

Experiment  65.  Detection  of  Nitric  Acid.  —  (Two  test-tubes ; 
nitric  acid ;  a  nitrate ;  concentrated  sulphuric  acid ;  ferrous 
sulphate.) 

Into  one  test-tube  introduce  the  nitric  acid  or  the  nitrate, 
and  some  concentrated  sulphuric  acid.  Warm  until  the  tube 
feels  warm  to  the  hand. 

Fill  the  other  test-tube  about  one-third  full  with  crystals  of 
green  ferrous  sulphate.  Add  just  enough  water  to  cover  the 
crystals  and  warm  until  the  tube  feels  warm  to  the  hand. 

Slant  both  tubes  and  carefully  pour  a  few  drops  from  one  into 
the  other.  Where  the  two  solutions  come  in  contact  a  dark 
purplish  ring  will  be  formed.  This  is  the  general  test  for  nitric 
acid.  (See  page  84.) 


CHAPTER  X 

NEUTRALIZATION   OF   ACIDS   AND   BASES 

Bases  are  Hydroxyl  Compounds.  —  Ammonium  hydroxide 
when  dissolved  in  water  dissociates  into  the  cation  ammonium, 
and  the  anion  hydroxyl.  The  hydroxyl  ion  and  not  the 
ammonium  ion  gives  the  characteristic  basic  property  to  the 
solution.  This  is  shown  by  the  fact  that  there  are  many  com- 
pounds which,  when  dissolved  in  water,  dissociate,  yielding  the 
ammonium  ion,  and  these  solutions  have  no  basic  properties. 
On  the  other  hand,  every  compound  which  yields  hydroxyl  ions 
is  a  basic  substance.  That  the  statement  that  bases  are 
hydroxyl  compounds  is  correct  can  be  seen  at  once  by  examin- 
ing the  composition  of  a  number  of  basic  substances :  — 

Sodium  hydroxide  ....  NaOH 

Potassium  hydroxide  ....  KOH 

Calcium  hydroxide  ....  Ca(OH)2 

Barium  hydroxide  ....  Ba(OH)2 

Aluminium  hydroxide  ....  A1(OH)3 

This  list  of  basic  substances  could  be  greatly  extended.  It 
will  be  observed  that  they  all  contain  a  metal  combined  with 
one  or  more  hydroxyl  groups.  When  these  substances  disso- 
ciate, the  hydroxyls  split  off  as  anions,  and  the  metal  forms 
the  cation.  A  few  examples  will  make  this  clear :  — 

NaOH     =  N+a,  OH. 
KOH       =  K,  OH. 
Ca(OH)2  =  Ca,  OH,  OH. 

Ba(OH)2  =  Ba,  OH,  OH. 

+++    —       —       — 
A1(OH)3  =  Al,  OH,  OH,  OH. 


NEUTRALIZATION   OF  ACIDS  AND  BASES  99 

Acidity  of  Bases  and  Basicity  of  Acids.  —  We  observe  in  the 
above  examples  that  some  bases  dissociate  yielding  one 
hydroxyl  ion,  other  bases  yield  two  hydroxyl  ions,  and  others 
still,  three  hydroxyl  ions.  If  we  take  a  gram-molecular  weight 
(the  molecular  weight  of  the  substance  in  grams)  of  a  mono- 
basic acid  and  dissolve  it  in  water,  diluting  the  solution  to  a 
litre,  and  take  a  gram-molecular  weight  of  any  one  of  the  above 
bases  which  yields  one  hydroxyl  ion  and  dissolve  it  so  as  to 
form  a  litre  of  solution,  the  litre  of  the  acid  would  exactly 
neutralize  the  litre  of  the  base.  Such  a  base  which  yields  on 
dissociation  one  hydroxyl  ion  is  known  as  a  monacid  base :  — 

H,  01  +  Na,  OH  =  H20  +  Na,  01. 

If  we  prepare  a  solution  of  a  base  which  dissociates  into  two 
hydroxyl  ions,  containing  a  gram-molecular  weight  in  a  litre,  it 
will  require  just  two  litres  of  a  solution  of  an  acid  such  as  that 
referred  to  above  to  neutralize  the  one  litre  of  the  base :  — 

Ca,  OH,  OH  +  H,  01  +  H,  01  =  2  H20  +  Ca,  01,  01. 

Such  bases  are  known  as  diacid  bases.  To  neutralize  a  gram- 
molecular  weight  of  a  base  which  dissociates  into  three  hy- 
droxyl ions,  requires  just  three  litres  of  the  above  solution  of 
acid.  Such  a  base  is  termed  a  triacid  base :  — 

+++    —       —      —          +—        +     —        -f     — 
Al,  OH,  OH,  OH,  +  H,  Cl  +  H,  01  +  H,  01  = 

3  H20  ++£L,  01,  01,  01. 

Just  as  we  have  mono,  di,  and  triacid  bases,  just  so  we  have 
mono,  di,  and  tribasic  acids.  An  acid  which  dissociates,  yield- 
ing one  hydrogen  ion,  is  monobasic :  — 

HC1  =  H,  01. 

If  the  molecule  of  the  acid  yields  two  hydrogen  ions,  it  is 
dibasic :  — 

H2S04  =  H,  H,  SO4. 


100  ELEMENTS  OF  INORGANIC  CHEMISTRY 

If  the  molecule  of  the  acid  dissociates,  yielding  three  hy- 
drogen ions,  it  is  tribasic :  — 

H3As04  =  H,  H,  H,  As04; 

and  so  on. 

Salts.  —  When  a  dilute  solution  of  an  acid  acts  on  a  dilute 
solution  of  a  base,  what  takes  place  and  all  that  takes  place  is 
the  formation  of  a  molecule  of  water :  — 

01,  H  +  OH,  Na  =  H20  +Na,  01. 

The  sodium  ion  remains  after  the  process  of  neutralization 
in  exactly  the  same  condition  as  before,  and,  similarly,  the 
chlorine  remains  in  the  ionic  condition.  The  hydrogen  and 
hydroxyl  ions,  however,  unite  and  form  a  molecule  of  water. 
It  is  a  general  rule  that,  whenever  we  have  hydrogen  and 
hydroxyl  ions  in  the  presence  of  one  another  uncombined,  they 
unite  and  form  water.  There  is  an  abundance  of  direct  experi- 
mental evidence  in  favor  of  this  conclusion. 

If,  however,  we  evaporate  the  solution  containing  the  sodium 
and  chlorine  ions,  they  unite  and  form  a  molecule  of  sodium 
chloride.  This  is  a  salt.  We  would  define  a  salt  as  follows : 
A  salt  is  a  compound  formed  by  the  union  of  an  anion  of  an  acid 
with  a  cation  of  a  base.  This  takes  place  generally,  as  already 
stated,  only  when  the  solution  containing  these  ions  is  evapo- 
rated, and  at  least  a  part  of  the  water  removed. 

The  salts  are  named  after  the  acids  from  which  they  are 
derived.  Salts  of  hydrochloric  acid  are  called  chlorides,  those 
of  nitric  acid  nitrates,  and  those  of  sulphuric  acid  sulphates. 
In  general,  salts  of  acids  which  end  in  "  ic "  are  termed 
"ates."  Salts  of  sulphurous  acid  are  called  sulphites,  salts 
of  nitrous  acids  nitrites,  and  so  on.  In  general,  the  names  of 
salts  of  acids  which  end  in  "  ous,"  end  in  "  ite." 

So  much  for  the  nomenclature  of  salts  in  terms  of  the  acids. 
Since  the  cation  also  enters  into  the  salt,  we  must  be  able  to 
distinguish  the  salts  of  one  cation  from  the  salts  of  another 
cation.  The  name  of  the  cation  is  used  before  the  name  of  the 


NEUTRALIZATION  OF  ACIDS  AND  BASES  101 

acid  with,  whose  salt  we  are  dealing.  Thus,  the  chloride  of 
sodium  is  known  as  sodium  chloride,  the  chloride  of  calcium, 
calcium  chloride,  and  so  on.  When  we  come  to  a  metal  that 
shows  different  valence  the  case  is  a  little  more  complicated. 

Heat  of  Neutralization.  —  When  solutions  of  acids  and  bases 
are  brought  together,  heat  is  liberated.  Quantitative  measure- 
ments of  the  amounts  of  heat  set  free  brought  out  a  simple 
and  very  important  relation.  This  can  best  be  seen  from  the 
following  results  for  strong  acids  and  bases.  Gram-molecular 
weights  of  different  acids  were  brought  together  with  a  gram- 
molecular  weight  of  a  given  base,  both  the  acid  and  base  being 
present  in  very  dilute  solution  :  — 

HEAT  OF 
NEUTRALIZATION 

Hydrochloric  acid  and  sodium  hydroxide  .  .  13,700  cals. 

Nitric  acid  and  sodium  hydroxide         .  .  .  13,700  cals. 

Chloric  acid  and  sodium  hydroxide       .  .  .  13,760  cals. 

Bromic  acid  and  sodium  hydroxide       .  .  .  13,780  cals. 

The  remarkable  fact  comes  out  that  the  heat  of  neutralization 
of  these  strong  acids  with  a  given  base,  sodium  hydroxide,  is  a 
constant. 

This  suggests  a  farther  question  very  closely  correlated  to 
the  above.  Suppose  we  neutralize  a  given  acid  with  a  number 
of  bases,  will  the  heat  liberated  be  a  constant  ?  and  if  so,  will 
this  bear  any  close  relation  to  the  above  constant  where  the 
base  was  the  same  and  the  acid  changed?  This  can  be 
answered  by  the  following  results :  — 

HEAT  OF 
NEUTRALIZATION 

Hydrochloric  acid  and  lithium  hydroxide     .  .  13,700  cals. 

Hydrochloric  acid  and  potassium  hydroxide  .  13,700  cals. 

Hydrochloric  acid  and  barium  hydroxide     .  .  13,800  cals. 

Hydrochloric  acid  and  calcium  hydroxide    .  .  13,900  cals. 

The  heat  of  neutralization  of  a  given  acid  with  a  number  of 
bases  is  also  a  constant.  But  what  is  even  more  surprising, 
the  constant  in  this  case  has  the  same  value  as  in  the  preceding 


102  ELEMENTS   OF   INORGANIC   CHEMISTRY 

case  where  the  base  was  unchanged,  and  the  nature  of  the  acid 
varied. 

These  facts  when  they  were  first  discovered  were  very  per- 
plexing. Indeed,  no  satisfactory  explanation  of  them  could  be 
furnished,  and  it  was  not  until  the  theory  of  electrolytic  disso- 
ciation was  proposed  that  we  could  account  for  them  at  all. 

Explanation  of  the  Constant  Heat  of  Neutralization  of  Strong 
Acids  and  Strong  Bases.  —  It  is  one  of  the  crowning  glories  of 
the  theory  of  electrolytic  dissociation,  that  it  not  only  explains 
all  of  the  facts  in  connection  with  the  neutralization  of  strong 
acids  and  bases  in  dilute  aqueous  solution ;  but  these  facts  are 
a  necessary  consequence  of  the  theory. 

Take,  as  an  example,  hydrochloric  acid  and  sodium  hy- 
droxide. In  a  very  dilute,  aqueous  solution  of  hydrochloric 
acid  all  the  molecules  are  dissociated  into  hydrogen  ions  and 
chlorine  ions  thus:  — 

HC1  =  H,  Cl. 

Similarly,  in  dilute,  aqueous  solution  the  molecules  of  sodium 
hydroxide  are  completely  broken  down  into  ions :  — 

NaOH  =  Na,  OH. 

When  the  dilute  aqueous  solutions  of  the  base  and  acid  are 
brought  together,  the  following  reaction  takes  place :  — 

Na,  OH  +  H,  Cl  =  Na,  Cl  +  H20. 

The  cation  of  the  base,  sodium,  and  the  anion  of  the  acid, 
chlorine,  remain  in  solution  as  ions  after  the  process  of  neutral- 
ization in  exactly  the  same  condition  as  before  neutralization 
took  place.  The  anion  of  the  base,  hydroxyl,  and  the  cation 
of  the  acid,  hydrogen,  combine  and  form  a  molecule  of  water. 
This  is  proved  by  the  fact  that  water  is  always  formed  as  the 
result  of  the  process  of  neutralization. 

Since  hydroxyl  is  the  anion  of  every  base,  and  hydrogen  the 


NEUTRALIZATION  OF  ACIDS   AND   BASES  103 

cation  of  every  acid,  the  process  of  neutralization  of  any  strong 
acid  with  any  strong  base  in  dilute  solution,  consists  in  the 
union  of  the  hydroxyl  ion  of  the  base  with  the  hydrogen  ion 
of  the  acid,  forming  a  molecule  of  water. 

The  process  of  neutralization  of  any  acid  by  any  base  is, 
therefore,  exactly  the  same  as  the  process  of  neutralization  of 
any  other  acid  by  any  other  base.  The  total  heat  that  is  liber- 
ated when  a  gram-equivalent  of  a  completely  dissociated  acid 
acts  on  a  gram-equivalent  of  a  completely  dissociated  base,  is 
the  heat  set  free  by  the  union  of  a  gram-equivalent  of  hydroxyl 
ions  with  a  gram-equivalent  of  hydrogen  ions.  Thus :  — 


H  aq  +  OH  aq  =  13,700  cals. 


Since  all  processes  of  neutralization  of  completely  dissociated 
acids  and  bases  are  the  same,  the  heat  of  neutralizatian  of  all 
such  acids  and  bases  must  be  a  constant. 

Indicators.  —  In  order  to  determine  when  an  acid  and  a  base 
just  neutralize  one  another,  certain  substances  are  used  which 
have  a  different  color  in  the  presence  of  an  acid  from  that  in 
the  presence  of  a  base.  Such  substances  are  termed  indicators. 
Take  the  organic  coloring  matter  litmus.  In  the  presence  of 
an  acid  it  is  red.  In  the  presence  of  a  base  it  is  blue.  In 
neutral  solution  it  is  purple.  By  adding  a  few  drops  of  a 
solution  of  litmus  to  any  given  solution  we  can  tell  whether 
the  solution  is  acidic,  basic,  or  neutral. 

If  we  wish  to  just  neutralize  a  base  with  an  acid,  we  add  the 
indicator  to  the  solution  of  the  base.  The  indicator  is  colored 
blue.  Add  the  acid  drop  by  drop  until  the  solution  becomes 
purple,  which  means  that  all  the  base  has  been  neutralized 
and  that  the  solution  is  neutral. 

If  we  use  the  weak  organic  acid  phenolphthalein,  we  will 
find  that  it  is  purplish  red  in  the  presence  of  bases,  and  color- 
less in  the  presence  of  acids. 

Methyl  orange,  on  the  other  hand,  is  deep-  red  in  the  pres- 
ence of  acids,  and  yellow  in  the  presence  of  alkalies. 


104  ELEMENTS   OF  INORGANIC  CHEMISTRY 


EXPERIMENTS  WITH  ACIDS  AND  BASES 

Experiment  66.  Neutralization  of  Acids  by  Bases.  —  (Hydro- 
chloric acid ;  sulphuric  acid ;  nitric  acid ;  sodium  hydroxide ; 
potassium  hydroxide  ;  barium  hydroxide.) 

Introduce  10  cc.  of  dilute  hydrochloric  acid  into  a  beaker 
containing  50  cc.  of  water.  Add  a  few  drops  of  litmus.  Just 
neutralize  the  solution  with  sodium  hydroxide.  If  the  water 
is  now  removed  by  evaporation,  the  cation  of  the  base  sodium 
will  combine  with  the  anioii  of  the  acid  chlorine,  and  the  salt, 
sodium  chloride,  will  be  formed.  This  will  appear  as  a  white 
mass  when  the  solution  is  evaporated  to  dryness. 

Proceed  in  an  exactly  similar  manner  with  potassium  hydrox- 
ide and  sulphuric  acid.  On  evaporating  the  water  the  cations  of 

potassium  will  unite  with  the  sulphuric  anion,  S04,  forming 
potassium  sulphate. 

Perform  the  same  experiment  with  nitric  acid  and  barium 
hydroxide.  On  evaporation  the  barium  nitrate  will  separate  as 
a  white  powder. 

In  this  connection  see  Experiment  44  in  which  salts  of 
ammonium  were'  formed.  (See  page  100.^) 

Experiment  67.  Heat  of  Neutralization.  —  (Dilute  hydro- 
chloric acid;  dilute  sodium  hydroxide.) 

Introduce  50  cc.  of  dilute  hydrochloric  acid  into  one  beaker, 
and  an  equal  volume  of  dilute  sodium  hydroxide  into  another 
beaker.  Introduce  a  thermometer  into  each  beaker  and  note 
the  temperatures.  Pour  slowly  the  acid  into  the  alkali,  with 
vigorous  stirring.  Bead  the  thermometer  immersed  in  the 
mixture.  What  do  you  note  ?  Explain.  (See  page  101.) 

Experiment  68.  Indicators.  —  (Dilute  acid  ;  dilute  alkali ; 
litmus  ;  phenolphthalein ;  methyl  orange.) 

Introduce  10  cc.  of  dilute  sodium  hydroxide  into  a  beaker 
containing  100  cc.  of  water,  and  add  a  few  drops  of  litmus  until 
the  solution  shows  a  distinctly  blue  color.  Add  acid  drop  by 
drop,  stirring  vigorously  with  addition  of  each  drop.  As  soon 
as  the  base  is  neutralized  and  even  a  slight  excess  of  acid  is 
added,  the  solution  will  turn  deep  red  in  color.  Now  add  alkali, 
and  the  blue  color  will  return  again.  As  we  have  already 
seen,  the  blue  color  of  litmus  is  characteristic  of  the  presence 
of  an  alkali,  and  the  red  color  shows  that  an  acid  is  present. 

Repeat  the  same  experiment,  adding  to  the  alkali  a  few  drops 
of  an  alcoholic  solution  of  phenolphthalein.  The  beautiful  red 


NEUTRALIZATION  OF  ACIDS  AND   BASES 


105 


color  is  a  characteristic  reaction  of  this  indicator  with  alkalies. 
Add  hydrochloric  acid  cautiously  from  a  pipette,  stirring  after 
the  addition  of  each  drop.  When  the  alkali  has  all  been  neutral- 
ized, the  red  color  will  disappear,  and  the  solution  will  become 
colorless.  Now  add  a  little  alkali,  and  the  red  color  will  return 
again. 

Repeat  the  same  experiment,  using  methyl  orange.  In  the 
presence  of  an  alkali  it  is  slightly  yellow  in  color.  In  the  pres- 
ence of  an  acid  it  is  deep  red. 

These  reactions  show  some  of  the  methods  which  are  at  our 
disposal  for  detecting  alkalies  and  acids.  (See  page  103.) 

Experiment  69.  Neutralization  of  Acids  and  Bases,  Quanti- 
tative Work.  —  (Two  burettes ;  dilute  hydrochloric  and  sulphu- 
ric acids ;  dilute  sodium  and  potassium  hydroxides ;  methyl 
orange ;  litmus.) 

Add  10  cc.  of  ordinary  dilute  hydrochloric  acid  to  a  500  cc. 
measuring-flask,  and  fill  the  flask  to  the  mouth  with  distilled 
water.  Shake  from  time  to  time. 
Pour  the  solution  into  a  bottle,  and 
shake  repeatedly  until  used. 

In  a  similar  manner  prepare  very 
dilute  solutions  of  sulphuric  acid, 
sodium  hydroxide,  and  potassium 
hydroxide,  using  5  cc.  of  "  dilute " 
sulphuric  acid  and  20  cc.  of  each  of 
the  "  dilute  "  bases,  and  diluting  each 
to  500  cc.  The  apparatus  to  be  used 
is  shown  in  Figure  33.  It  consists  of 
two  burettes  supported  on  a  conven- 
ient stand.  A  burette  is  a  long,  nar- 
row, glass  tube,  graduated  to  cubic 
centimetres.  It  is  closed  at  the  bot- 
tom by  a  rubber  tube  grasped  by  a 
pinch-cock.  By  pressing  on  the  pinch- 
cock  the  tube  is  opened,  and  the  liquid 
can  flow  out  of  the  tube. 

Fill  one  of  the  burettes  with  the 
dilute  solution  of  hydrochloric  acid, 
and  the  other  with  the  dilute  solution 
of  sodium  hydroxide.  Run  25  cc.  of 
the  alkali  into  each  of  three  small 
beakers  and  add  two  or  three  drops  of  methyl  orange  to  each 
beaker.  Run  in  the  acid  drop  by  drop,  stirring  vigorously 


FIG.  33. 


106  ELEMENTS   OF  INORGANIC   CHEMISTRY 

after  the  addition  of  each  drop.  Watch  carefully  for  the  ap- 
pearance of  the  first  trace  of  the  red  color  of  the  indicator. 
This  means  that  all  of  the  alkali  has  been  neutralized.  Having 
read  the  burette  before  any  acid  was  admitted,  read  it  again. 
Note  the  amount  of  acid  required  to  neutralize  25  cc.  of  the 
alkali.  Repeat  the  experiment  three  times  and  see  how  closely 
the  results  agree  in  the  three  determinations. 

Repeat  the  same  experiment,  using  potassium  hydroxide 
instead  of  sodium  hydroxide. 

Repeat  the  above  experiments,  using  hydrochloric  acid  with 
both  bases,  and  note  the  results. 

Repeat  one  or  two  of  the  last  experiments,  using  litmus  as 
the  indicator  instead  of  methyl  orange.  Study  thoroughly  the 
numerical  data  obtained,  and  it  will  be  found  that  a  given 
amount  of  acid  always  requires  a  given  amount  of  base  to 
neutralize  it.  Show  that  this  is  the  case. 


CHAPTER   XI 

THE  ATMOSPHERIC  AIR   AND   CERTAIN   RARE 
ELEMENTS   OCCURRING  IN  IT 

THE   ATMOSPHERIC   AIR 

IT  was  stated  when  we  were  studying  nitrogen  that  the  chief 
source  of  that  element  was  the  atmospheric  air.  Indeed,  it 
comprises  nearly  four-fifths  of  the  atmosphere.  In  addition  to 
nitrogen,  we  find  an  abundance  of  oxygen  in  the  atmosphere. 
This  amounts  to  nearly  one-fifth  of  the  whole.  In  addition 
to 'these  two  elements  we  find  many  other  substances,  both 
elementary  and  compound,  in  the  atmosphere,  so  that  we  must 
study  this  mixture  of  gases  with  some  thoroughness. 

Composition  of  the  Atmosphere.  —  In  order  to  determine  the 
exact  composition  of  the  atmosphere,  we  must  make  a  quantita- 
tive analysis  of  it.  The  oxygen  in  the  air  can  be  determined 
in  several  ways.  A  measured  volume  of  air  can  be  passed  over 
heated  copper.  The  oxygen  combines  with  the  copper,  forming 
copper  oxide.  By  weighing  the  tube  containing  the  copper 
before  the  experiment,  and  weighing  the  tube  containing  the 
copper  and  copper  oxide  after  the  experiment,  we  know  from 
the  gain  in  weight  the  weight  of  the  oxygen  in  a  given  vol- 
ume of  air.  Knowing  the  weight  of  a  litre  of  oxygen  and  of 
a  litre  of  air,  the  percentage  of  oxygen  can  be  calculated  at 
once. 

Again,  the  oxygen  can  be  removed  from  the  air  by  inserting 
a  piece  of  phosphorus,  which  will  combine  with  the  oxygen. 
By  measuring  the  original  volume  of  the  air,  and  the  volume 
after  all  the  oxygen  has  been  removed,  we  have  the  percentage  of 
oxygen  by  volume  in  the  atmospheric  air. 

107 


108  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Results  by  the  different  methods  show  that  the  oxygen  in 
the  air  is  about  20.8  per  cent  by  volume,  and  23  per  cent  by 
weight. 

The  question  arises,  Does  the  amount  of  oxygen  present 
remain  constant,  or  does  it  vary  from  place  to  place  or  from 
time  to  time  ?  While  slight  variations  have  been  detected,  pure 
air  from  different  parts  of  the  globe,  and  in  different  altitudes, 
varies  very  slightly  in  composition. 

The  remainder  of  the  atmospheric  air  is  nearly  all  nitrogen,  a 
number  of  other  substances,  however,  occurring  in  it  in  very 
small  quantities.  There  are  traces  of  carbon  dioxide  in  the 
air.  The  amount  can  be  determined  by  passing  the  air  through 
a  solution  of  barium  hydroxide,  and  weighing  the  amount  of 
barium  carbonate  precipitated. 

The  small  quantity  of  ammonia  in  the  air  can  be  determined 
by  passing  a  given  volume  of  air  through  a  solution  of  a  stand- 
ard acid,  and  determining  how  much  of  the  acid  is  neutralized. 

The  air  under  all  conditions  contains  water-vapor.  The 
amount,  however,  varies  greatly  from  time  to  time  and  from 
place  to  place. 

To  determine  the  amount  of  water-vapor  in  the  atmosphere, 
it  is  only  necessary  to  pass  a  measured  volume  of  air  over  some 
good  drying  agent,  such  as  phosphorus  pentoxide,  and  deter- 
mine the  increase  in  the  weight  of  the  pentoxide. 

Other  substances  may  occur  in  the  atmosphere  in  very  minute 
quantities,  such  as  ozone,  hydrogen  dioxide,  oxides  of  nitrogen, 
and  the  like,  but  the  quantities  are  so  small  that  they  can,  for 
all  practical  purposes,  be  disregarded. 

In  addition  to  the  constituents  already  named,  there  are  a 
number  of  rare  elements  which  occur  in  the  air  in  very  small 
quantities.  These  are  the  newly  discovered  elements,  argon, 
helium,  neon,  krypton,  and  xenon.  These  elements  we  shall 
consider  briefly  a  little  later. 

Is  the  Air  a  Mixture  or  a  Compound?  —  The  question  natu- 
rally arises,  Is  the  atmospheric  air  a  chemical  compound  or  a 
mechanical  mixture  ?  The  fact  that  it  has  so  nearly  the  same 


7 

THE   ATMOSPHERIC  AIR  109 

composition  the  world  over,  would  argue  in  favor  of  the  oxygen 
and  nitrogen  being  in  combination,  forming  a  definite  com- 
pound. This  line  of  argument,  however,  is  by  no  means  con- 
clusive, since  we  might  easily  have  the  two  gases  mixed  in 
essentially  the  same  proportion  in  all  regions.  There  is, 
however,  direct  evidence  which  shows  that  the  air  is  simply 
a  mechanical  mixture  of  oxygen  and  nitrogen,  and  not  a 
chemical  compound. 

When  air  is  shaken  with  water,  the  part  that  dissolves  has 
a  very  different  composition  from  ordinary  air.  The  latter 
contains  in  round  numbers  four  parts  of  nitrogen  to  one  of 
oxygen,  while  air  which  has  been  dissolved  in  water  contains 
only  1.9  parts  of  nitrogen  to  one  of  oxygen.  This  is  due  to  the 
fact  that  oxygen  is  much  more  readily  soluble  in  water  than 
nitrogen.  If  air  is  a  compound  of  oxygen  and  nitrogen,  the 
compound  would  dissolve  as  such,  and  the  air  which  would  be 
dissolved  by  water  would  have  the  same  composition  as  ordinary 
air. 

Again,  chemical  union  is  always  accompanied,  as  we  express 
it,  by  thermal  change.  Oxygen  and  nitrogen  mix  in  the  pro- 
portion to  form  air  without  any  thermal  change,  and  air  is, 
therefore,  not  a  chemical  compound. 

Physical  Properties  of  Atmospheric  Air.  —  The  specific 
gravity  of  air  varies  slightly,  just  as  the  composition  changes 
slightly.  Under  the  average  conditions  of  0°  and  760  mm. 
pressure,  one  litre  of  air  weighs  1.293  grams.  The  pressure 
of  the  air,  however,  decreases  very  rapidly  as  we  rise  from  the 
level  of  the  sea,  and  a  litre  of  air  on  the  top  of  a  high  mountain 
would  weigh  much  less. 

^  Liquid  Air.  —  We  have  seen  that  both  oxygen  and  nitrogen 
can  be  liquefied,  and  would  expect,  therefore,  that  atmospheric 
air,  which  is  essentially  a  mixture  of  these  two  gases,  could 
also  be  liquefied.  Such  is  the  fact.  The  method  employed  is 
based  on  exactly  the  same  principles  that  were  made  use  of  to 
liquefy  oxygen  and  similar  substances.  The  most  economical 
method  consists  in  compressing  the  air  and  removing  the  heat 


110  ELEMENTS   OF  INORGANIC   CHEMISTRY 

set  free  by  a  stream  of  cold  water.  The  compressed  air  is  allowed 
to  expand,  when  its  temperature  is  very  much  lowered.  It  is 
then  allowed  to  cool  other  compressed  air,  which,  in  turn,  is 
allowed  to  expand,  and  a  still  lower  temperature  is  produced. 
This  is  continued  until  a  temperature  is  reached  at  which  the 
compressed  air,  when  allowed  to  expand,  becomes  partly  lique- 
fied. In  this  process  the  air  is  allowed  to  expand  through  a 
fine  opening  known  as  a  needle  valve,  when  part  of  the  com- 
pressed air  is  liquefied  and  the  remainder  passes  off  as -gas. 

Liquid  air  has  a  slightly  bluish  color.  When  filtered  from 
solid  carbon  dioxide  and  ice  it  is  transparent.  It  boils  at 
—  190°.  The  liquid  nitrogen,  having  a  lower  boiling-point  than 
liquid  oxygen,  boils  off  more  rapidly,  and  the  liquid  remaining 
after  liquid  air  has  been  allowed  to  evaporate  for  a  considerable 
time,  is  almost  pure  liquid  oxygen.  \7 

ARGON,  HELIUM,  KRYPTON,  NEON,  XENON 

Argon.  —  These  five  elements  have  all  been  discovered  in 
atmospheric  air  since  the  summer  of  1894.  Just  before  this 
time  Lord  E-ayleigh  had  observed  that  nitrogen  obtained  from 
atmospheric  air  by  removing  all  known  constituents  was  slightly 
heavier,  volume  for  volume,  than  nitrogen  prepared  by  heating 
ammonium  nitrite. 

The  most  probable  explanation  seemed  to  be  that  the  nitrogen 
from  the  air  contained  some  impurity  which  was  heavier  than 
nitrogen. 

Kayleigh  and  Eamsay  each  isolated  the  gas,  using,  however, 
very  different  methods.  Ramsay  proceeded  as  follows :  He 
determined  to  remove  the  oxygen  from  the  air,  then  the 
nitrogen  and  other  known  constituents,  and  see  if  anything 
remained. 

He  removed  the  oxygen  from  the  air  by  passing  it  over  red- 
hot  copper.  The  nitrogen  was  removed  from  the  residue  by 
passing  it  over  red-hot  magnesium,  the  ordinary  impurities 
having  been  previously  removed.  There  remained  a  residue 


THE   ATMOSPHERIC   AIR  111 

which  spectrum  analysis  showed  to  be  a  new  substance,  and 
which  was  a  little  less  than  one  per  cent  of  the  atmosphere. 
Ramsay  and  Eayleigh  were  not  able  to  make  it  combine  with 
any  known  substance,  and  from  its  chemical  inactivity  called 
it  argon. 

Helium,  Neon,  Krypton,  and  Xenon.  —  Since  the  discovery  of 
argon,  Ramsay  has  carried  his  investigations  on  the  atmospheric 
air  much  farther,  and  has  discovered  four  new  substances, 
all  of  which  appear  to  be  elementary.  When  air  is  liquefied, 
two  of  these  escape,  being  very  volatile  —  helium  and  neon. 

Helium,  so  called  because  it  had  been  recognized  by  means 
of  the  spectroscope  as  occurring  in  the  sun,  has  also  been  dis- 
covered in  the  waters  of  certain  springs,  and  in  certain  ores  of 
uranium.  When  a  mixture  of  helium  and  neon  is  cooled  in 
liquid  hydrogen,  the  neon  is  liquefied,  while  the  helium  remains 
a  gas.  Helium  does  not  combine  with  any  known  substance, 
and  its  boiling-point  is  somewhat  lower  than  that  of  hydrogen. 
It  has,  then,  the  lowest  boiling-point  of  any  known  substance, 
and  has  thus  far  not  been  liquefied. 

Krypton  and  Xenon  boil  higher  than  air,  and  were,  therefore, 
found  in  the  residue  from  the  evaporation  of  a  large  amount 
of  liquid  air.  They  were  separated  by  means  of  the  difference 
in  their  boiling-points. 

EXPERIMENTS  WITH  ATMOSPHERIC   AIR 

Experiment  70.  Determination  of  the  Amount  of  Oxygen 
in  the  Atmosphere.  —  (Tall  glass  cylinder;  glass  tube  closed 
at  one  end  holding  about  100  cc.,  and  graduated  into  centi- 
metres and  millimetres ;  a  piece  of  phosphorus  attached  to  a 
bent  wire.) 

Arrange  the  apparatus  as  in  Figure  1.  Before  introducing 
the  phosphorus,  bring  the  water  inside  the  tube  to  the  same 
level  as  outside  and  read  the  volume  of  the  inclosed  gas.  Read 
the  temperature  of  a  thermometer  attached  to  the  glass 
tube ;  also  the  height  of  the  barometer  in  the  same  room  in 
which  the  experiment  is  being  carried  out.  The  reasons  for 
making  these  readings  will  appear  a  little  later. 


112  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Introduce  the  phosphorus  and  allow  it  to  remain  in  the  tube 
for  at  least  12  hours.  Then  remove  the  phosphorus.  It  will 
be  seen  at  once  that  the  volume  of  the  gas  has  diminished. 
Bring  the  water  to  the  same  level  inside  and  outside  the  tube 
by  lowering  the  tube  in  the  cylinder ;  read  the  volume  of  the 
residual  gas,  also  the  temperature  and  barometric  pressure. 

The  reason  for  observing  the  temperature  and  pressure  when 
making  the  above  readings  will  appear  from  the  following  con- 
siderations. 

The  volume  of  a  gas  changes  with  the  pressure  to  which  the 
gas  is  subjected.  The  greater  the  pressure,  the  less  the  volume. 
Indeed,  we  know  from  Boyle's  law  for  gases  that  the  volume 
varies  inversely  as  the  pressure.  From  day  to  day  the  baro- 
metric pressure  changes  and  therefore  must  be  taken  into 
account. 

The  volume  of  a  gas  also  varies  with  the  temperature  —  the 
higher  the  temperature  the  greater  the  volume,  pressure  being 
constant.  For  every  rise  in  temperature  of  one  degree  centi- 
grade the  volume  of  a  gas  will  increase  =  0.003665,  which 

273 

is  the  law  of  Gay  Lussac  for  gases.  Since  temperature  changes 
from  time  to  time,  this  also  must  be  taken  into  account. 

In  order  that  the  volumes  of  gases  as  measured  may  be  com- 
parable, we  must  refer  them  to  comparable  conditions.  The 
conditions  which  have  been  selected  are  a  temperature  of  0°  C. 
and  a  pressure  of  760  mm.  of  mercury. 

Let  the  volume  of  the  gas  as  read  at  t°  and  p  mm.  of  mer- 
cury pressure  be  v.  The  volume  of  this  gas  "Fat  0°  and  760 
mm.  pressure  will,  from  the  above  relations,  be :  — 


760(1  +  0.003665  *) 

This  takes  into  account  every  factor  except  the  tension  of 
the  aqueous  vapor  in  the  tube.  This  pressure  acts  against  the 
barometric  pressure,  so  that  the  gas  is  really  under  less  pressure 
than  would  be  indicated  by  the  barometer.  This  must  be  sub- 
tracted from,  the  barometric  pressure ;  we  will  represent  the 
tension  of  the  aqueous  vapor  by  h.  The  true  volume  of  the 
gas  Fat  0°  and  760  mm.  pressure  would  then  be  :  — 

v=          v(p-Ji) 

760(1  +  0.003665*) 


THE  ATMOSPHERIC  AIR 


113 


In  both  of  the  above  readings  calculate  the  true  volume  of 
the  gas  under  standard  conditions.  It  will  be  found  that 
about  one-fifth  of  the  gas  has  combined  with  the  phosphorus, 
or,  more  accurately,  21  per  cent. 

Air,  therefore,  contains  approximately  21  per  cent  of  oxygen. 
That  the  remaining  gas  is  not  oxygen  can  be  shown  by  plung- 
ing into  it  a  lighted  match.  It  is  mainly  nitrogen,  but  contains 
quite  a  large  number  of  other  elements  and  compounds  as  we 
have  seen ;  but  these  are  present  in  very  small  quantities. 
(See  page  107.) 

Experiment  71.  Detection  of  Carbon  Dioxide  in  the  Atmos- 
phere. —  (A  large  bottle  ;  two  glass  cylinders ;  corks ;  glass 
tubing ;  rubber  tubing ;  pinch-cock  j  barium  hydroxide.) 


FIG  34. 

Arrange  an  apparatus  as  shown  in  Figure  34.  The  large  glass 
bottle  C  is  filled  with  water.  The  glass  cylinders  contain  a 
solution  of  barium  hydroxide.  Allow  the  water  to  flow  slowly 
out  of  the  end  of  the  tube,  by  opening  the  pinch-cock,  which 
closes  the  rubber  tube.  As  the  water  flows  out,  air  will  pass  in 
through  the  cylinders  A  and  B,  into  C.  A  white  precipitate 
will  be  formed  in  A  and  B,  but  will  be  much  heavier  in  A  than 
in  B.  Barium  hydroxide  combines  readily  with  carbon  dioxide, 
forming  barium  carbonate,  which  is  a  white  solid,  insoluble  in 
water.  Most  of  the  carbon  dioxide  in  the  air  combines  with 


114  ELEMENTS  OF   INORGANIC   CHEMISTRY 

the  barium  hydroxide  in  the  first  cylinder;  a  little,  however, 
passes  on  to  the  second  cylinder  and  there  combines  with 
barium  hydroxide.  This  explains  why  a  heavier  precipitate 
is  formed  in  A  than  in  B. 

After  the  vessel  C  has  become  half  full  of  air,  remove  the 
cork  and  add  quickly  some  barium  hydroxide  —  closing  the  ap- 
paratus again  as  quickly  as  possible.  The  barium  hydroxide 
in  C  will  become  only  slightly  cloudy,  showing  that  practically 
all  of  the  carbon  dioxide  has  been  removed  from  the  air  in 
passing  through  the  solutions  in  A  and  B.  (See  page  108.) 

Experiment  72.  Proof  of  the  Presence  of  Water  in  the  Air.  — 
(Two  large  ground-glass  stoppered  weighing-tubes;  calcium 
chloride;  phosphorus  pentoxide.) 

Fill  a  ground-glass  stoppered  weighing-tube  about  one-half 
full  with  calcium  chloride.  Insert  the  stopper  and  weigh  the 
tube.  Remove  the  stopper  and  expose  to  the  air  for  several 
hours.  What  evidence  have  you  that  water  has  been  absorbed  ? 
Insert  the  stopper  and  reweigh  the  tube.  Has  it  gained  in 
weight  ? 

Perform  exactly  the  same  experiment  with  phosphorus  pen- 
toxide. Does  it  increase  in  weight  ?  (See  page  108.) 


CHAPTER  XII 

DETERMINATION   OF   RELATIVE  ATOMIC   "WEIGHTS 

Combining  Numbers  and  Atomic  Weights.  —  We  saw  in 
Chapter  VI  that  the  atomic  theory  was  proposed  to  account  for 
certain  well-established  laws  of  chemical  combination  —  the 
laws  of  definite  and  multiple  proportions.  If  atoms  are  the 
ultimate  units  of  matter,  these  must  have  definite  weights, 
and  it  is  obviously  of  great  importance  for  chemistry  to  deter- 
mine the  relative  weights  of  the  atoms  of  different  substances. 

If  the  same  number  of  atoms  of  any  two  substances  coin- 
bine,  the  relative  weights  of  the  substances  that  combine 
represent  the  relative  weights  of  the  atoms  which  enter  into 
combination.  This  apparently  furnishes  a  means  of  determin- 
ing relative  atomic  weights.  It  is  only  necessary  to  determine 
the  relative  weights  of  substances  that  combine  —  the  combin- 
ing numbers  —  in  order  to  ascertain  the  relative  weights  of 
the  atoms  of  these  substances.  This  would  be  true  if  a  given 
number  of  atoms  of  one  substance  always  combined  with  an 
equal  number  of  atoms  of  another.  But  we  know  that  this  is 
not  the  case,  since  it  often  happens  that  two  elementary  sub- 
stances combine  in  several  proportions.  To  determine  the 
relative  atomic  weights  of  the  elements,  we  must,  therefore, 
know  the  combining  numbers  of  the  elements,  and  also  the 
number  of  atoms  of  the  different  elements  which  combine  with 
one  another.  We  will  take  up  first  the  method  of  determining 
the  combining  numbers  of  the  elements. 

Chemical  Methods  of  determining  Combining  Numbers.  —  The 
simplest  method  would  be  to  take  some  element  as  our  stand- 
ard, and  call  its  combining  number  one.  Then  allow  all  of 

115 


116  ELEMENTS   OF   INORGANIC   CHEMISTRY 

the  other  elements  to  combine  with  this  one,  and  determine 
the  weights  of  the  different  elements  which  combined  with 
unit  weight  of  our  standard  element.  Since  hydrogen  has  the 
smallest  combining  number,  it  would  naturally  be  chosen  as 
the  unit.  The  problem  then  would  be  to  determine,  say,  the 
number  of  grams  of  the  different  elements  which  combine  with 
one  gram  of  hydrogen,  and  these  figures  would  represent  the 
combining  weights  of  the  elements  in  terms  of  hydrogen  as 
unity.  We  might  thus  work  out  a  table  of  the  combining 
numbers  of  all  of  the  elements  in  terms  of  hydrogen  as 
unity. 

This  part  of  the  problem  is,  however,  not  as  simple  as  would 
be  indicated  from  the  above.  Some  of  the  elements  combine 
with  hydrogen  in  more  than  one  proportion.  We  would  thus 
obtain  different  combining  numbers  for  the  same  element, 
depending  upon  which  of  its  compounds  we  selected. 

It  is  perfectly  clear  that  neither  the  chemical  analysis  of 
the  compound,  nor  its  synthesis  from  the  elements,  throws 
any  light  on  the  problem  as  to  the  number  of  atoms  of 
one  substance  combined  with  one  atom  of  the  other.  Other 
methods  must  be  employed  in  order  to  determine  the  number 
of  atoms  of  the  one  element  which  have  combined  with  one 
atom  of  the  other.  To  these  we  shall  now  turn. 

Avogadro's  Hypothesis.  —  Avogadro,  in  1811,  taking  into  ac- 
count all  of  the  facts  known,  advanced  the  hypothesis  that  — 

In  equal  volumes  of  all  gases,  at  the  same  temperature  and 
pressure,  there  is  an  equal  number  of  ultimate  parts  or  molecules. 

Avogadro  extended  his  hypothesis  to  all  gases,  including 
even  the  elementary  gases,  and  regarded  the  molecules  of 
these  substances  as  made  up  of  atoms  of  the  same  kind,  which 
had  united  with  one  another.  This  was  a  necessary  conse- 
quence of  his  hypothesis.  One  volume  of  hydrogen  gas  com- 
bines with  one  volume  of  chlorine  gas,  and  forms  two  volumes 
of  hydrochloric  acid  gas.  If  there  are  the  same  number  of 
molecules  in  equal  volumes  of  all  gases,  there  would  be  twice 
as  many  in  the  two  volumes  of  hydrochloric  acid  as  in  the  one 


DETERMINATION  OF  RELATIVE   ATOMIC   WEIGHTS      117 

volume  of  hydrogen,  or  the  one  volume  of  chlorine.  Since 
each  molecule  of  hydrochloric  acid  must  contain  at  least  one 
atom  of  hydrogen  and  one  atom  of  chlorine,  the  molecule  of 
hydrogen  and  of  chlorine  must  be  made  up  of  at  least  two 
atoms.  The  hypothesis  of  Avogadro  has  been  confirmed  by 
such  an  abundance  of  subsequent  work,  in  so  many  directions, 
that  it  is  now  placed  among  the  well-established  laws  of  nature. 

Avogadro's  Hypothesis  and  Molecular  Weights.  — Given  the 
hypothesis  of  Avogadro,  the  determination  of  the  relative 
molecular  weights  of  gases  is  very  simple.  If  there  is  an  equal 
number  of  molecules  contained  in  equal  volumes  of  the  dif- 
ferent gases,  the  relative  weights  of  equal  volumes  of  these 
gases  give  at  once  the  relative  weights  of  the  molecules  con- 
tained in  them.  It  is  only  necessary  to  choose  some  substance 
as  our  standard,  and  express  the  molecular  weights  of  other 
substances  in  terms  of  this  standard.  We  would  naturally 
select  as  the  unit  that  substance  which  has  the  smallest  density, 
and  this  is  hydrogen.  From  what  has  been  said,  however, 
in  reference  to  the  union  of  hydrogen  and  chlorine,  forming 
hydrochloric  acid,  it  is  certain  that  the  molecule  of  hydrogen 
contains  at  least  two  atoms.  We  will,  therefore,  call  the 
molecular  weight  of  hydrogen  two,  and  calculate  the  molecu- 
lar weights  of  other  elements  in  terms  of  this  standard.  The 
densities  of  substances  are  usually  determined  in  terms  of  air 
as  the  unit.  It  is  a  simple  matter  to  recalculate  these  in  terms 
of  hydrogen  as  two.  The  density  of  hydrogen  in  terms  of  air 
as  the  unit  is  0.0696.  We  must  multiply  this  by  28.73  to 
obtain  our  new  unit  2  (2  H-  0.0696  =  28.73).  Similarly,  for 
other  substances  whose  densities  are  known  with  reference 
to  air;  these  densities  must  be  multiplied  by  the  constant 
28.73  to  transform  them  into  densities  in  terms  of  hydrogen 
=  2.  These  latter  values  are  the  relative  molecular  weights 
of  the  substances  in  the  form  of  gas,  referred  to  the  molecular 
weight  of  hydrogen  as  two. 

Method  of  Dumas  for  determining  Vapor-densities.  —  The 
method  of  Dumas  consists  in  determining  the  amount  of  sub- 


118  ELEMENTS  OF  INORGANIC   CHEMISTRY 

stance  which,  in  the  form  of  vapor,  at  a  given  temperature,  just 
fills  a  flask  whose  volume  is  afterwards  ascertained.  The  flask 
is  weighed  full  of  air.  Knowing  the  volume  of  the  flask,  we 
know  the  weight  of  air  contained  in  it;  therefore,  we  know 
the  weight  of  the  empty  flask.  The  weight  of  the  flask  being 
known,  and  the  weight  of  the  flask  plus  the  substance  which 
just  filled  it  with  vapor,  we  know  the  weight  of  the  substance. 
By  determining  the  weights  of  the  vapors  of  different  sub- 
stances which  fill  a  flask  of 
given  volume,  we  have  the  rela- 
tive densities  of  the  vapors. 

The  apparatus  used  is  a  bal- 
loon flask  (Fig.  35)  holding 
from  200  to  300  cc. 

The  flask  is  carefully  dried 
and  weighed. 

A  few  grams  of  the  sub- 
stance whose  vapor-density  is 
to  be  determined  are  intro- 
duced into  the  flask,  the  neck 

drawn  out  to  a  capillary,  and 
FIG.  35. 

the    flask    placed    in    a   bath 

which  is  at  least  ten  or  fifteen  degrees  above  the  boiling-point 
of  the  substance.  The  substance  vaporizes,  and  drives  out  the 
air,  and  when  the  vapor  of  the  substance  ceases  to  escape,  the 
capillary  is  fused  shut.  The  flask  after  cooling  is  weighed. 
The  fine  point  is  then  cut  off  under  mercury  and  the  flask 
filled  with  mercury.  The  flask  may  then  be  weighed  again, 
or  the  mercury  poured  out  and  measured,  giving  the  volume  of 
the  flask. 

Results  of  Vapor-density  Measurements.  —  A  few  results  are 
given  in  the  following  table,  showing  in  column  I  the  densities 
in  terms  of  air  as  the  unit ;  in  column  II  the  densities  or  rela- 
tive molecular  weights  in  terms  of  hydrogen  =  2.  The  results 
in  column  II  are  obtained  by  multiplying  the  results  in  column 
I  by  28.73. 


DETERMINATION  OF   RELATIVE   ATOMIC   WEIGHTS      119 


i 

n 

Hydrogen,  0°  C. 

0.06926 

2 

Oxygen,  0°  C.     . 

1.10563 

31.76 

Nitrogen,  0°  C.    . 

0.9713 

27.90 

Sulphur,  1400°  C. 

2.17 

X  28.73 

62.34 

Chlorine,  200°  C. 

2.45 

70.38 

Bromine,  100°  C. 

5.54 

159.16 

Mercury,  1400°  C. 

6.81 

195.65 

Iodine,  940°  C.    . 

8.72 

250.52 

Atomic  Weights  from  Molecular  Weights.  —  If  we  knew  the 
number  of  the  atoms  contained  in  the  molecules  of  elements  in 
the  gaseous  state,  the  problem  of  relative  atomic  weights  would 
be  solved  at  once  by  dividing  the  molecular  weight  of  the  gas 
by  the  number  of  atoms  in  the  molecule.  The  problem  is,  how- 
ever, not  as  simple  as  this,  since  we  do  not  know  at  once  the 
number  of  atoms  in  the  molecules  of  elements.  Other  lines  of 
thought  have  enabled  us  to  solve  this,  the  second  part  of  our 
problem. 

The  definition  of  an  atom  as  an  indivisible  particle  of  matter 
shows  that  fractions  of  atoms  cannot  exist.  No  molecule  can 
contain  a  fraction  of  any  atom.  The  quantity  of  any  substance 
which  enters  into  a  molecule  must  be  at  least  one  atom.  It 
may  be  more  than  one,  but  it  cannot  be  less.  This  is  the  key 
to  the  problem.  Suppose  we  wish  to  determine  the  number  of 
hydrogen  atoms  in  a  molecule  of  hydrogen.  We  must  examine 
compounds  into  which  hydrogen  enters,  and  find  out  what  is 
the  smallest  quantity  of  hydrogen  which  enters  into  the  mole- 
cule of  the  compound.  Let  us  take  hydrochloric  acid,  whose 
molecular  weight  is  36.45.  This  is  shown  by  analysis  to  be 
composed  of  1  part  of  hydrogen  and  35.45  parts  of  chlorine. 
This  1  part  of  hydrogen  is  at  least  one  atom ;  it  may  be  more, 
but  it  cannot  be  less.  By  examining  a  large  number  of  com- 
pounds into  which  hydrogen  enters,  it  has  been  found  that 
hydrogen  never  enters  into  a  molecule  of  any  substance  in  a 


120 


ELEMENTS   OF  INORGANIC   CHEMISTRY 


smaller  quantity  than  in  hydrochloric  acid.  This  is,  therefore, 
for  us  the  atom  of  hydrogen,  but  it  may  in  reality  be  composed 
of  a  great  number  of  smaller  parts.  The  hydrogen  that  enters 
into  the  molecule  of  hydrochloric  acid  is  just  half  the  quantity 
that  forms  the  molecule  of  hydrogen  gas,  since  one  volume  of 
hydrogen  combining  with  one  volume  of  chlorine  yields  two 
volumes  of  hydrochloric  acid  gas.  The  molecule  of  hydrogen, 
therefore,  contains  at  least  two  atoms,  and  since  there  is  no 
experimental  reason  for  assuming  that  it  contains  more  than 
two,  we  say  that  the  molecule  of  hydrogen  is  made  up  by  the 
union  of  two  hydrogen  atoms.  Knowing  the  number  of  atoms 
in  the  molecule,  the  atomic  weight  follows  at  once  from  the 
molecular  weight  determined  by  vapor-density,  and  corrected  by 
the  most  refined  methods  of  chemical  analysis. 

By  methods  similar  to  the  above  the  molecules  of  many  ele- 
ments have  been  shown  to  be  composed  of  two  atoms.  But 
this  by  no  means  applies  to  all  elementary  substances.  The 
molecules  of  some  elementary  substances  contain  more  than 
two  atoms,  and  in  a  very  few  cases  the  molecule  and  atom  -seem 
to  be  identical. 

The  relations  between  the  molecular  weights  of  a  few  of  the 
elements  and  their  atomic  weights  are  given  in  the  following 
table :  — 


ELEMENTS 

ATOMIC  WEIGHTS 

MOLECULAB  WEIGHTS 

Hydrogen       .... 

1 

2 

Nitrogen         .... 

14.01 

28.02 

15.88 

31.76 

Phosphorus    .... 

30.96 

123.84 

Sulphur  ..... 

31.98 

f     63.96  above  800°  C. 

\    191.88  at  500°  C. 

Chlorine          .... 

35.18 

70.36 

Arsenic  ..... 

74.9 

299.6 

Cadmium        .... 

111.7 

111.7 

Mercury          .... 

199.8 

199.8 

DETERMINATION   OF    RELATIVE   ATOMIC   WEIGHTS      121 


This  table  brings  out  a  number  of  facts  of  interest.  The 
molecular  weight  of  a  number  of  the  elements  is  twice  as  great 
as  the  atomic  weight.  In  some  cases,  as  with  sulphur,  the 
molecular  weight  is  twice  the  atomic  weight  at  a  given  tem- 
perature, and  then  varies  with  the  temperature.  In  the  cases 
of  cadmium  and  mercury  the  molecular  weights  are  apparently 
identical  with  the  atomic  weights.  The  molecular  weights  of 
phosphorus  and  arsenic  are  four  times  the  atomic  weights. 

The  Relative  Weights  of  the  Atoms.  —  In  determining  atomic 
weights  we  must  choose  some  element  as  our  standard.  We 
would  naturally  take  the  lightest  element,  hydrogen,  and  call 
it  unity.  This  has  "been  done,  and  all  atomic  weights  referred 
to  this  unit.  But  it  is  unfortunately  true  that  hydrogen  does 
not  combine  directly  with  many  of  the  elements  and  form  stable 
compounds  that  can  be  analyzed. 

Oxygen,  on  the  other  hand,  does  combine  with  a  large  num- 
ber of  the  elements,  forming  some  of  the  most  stable  compounds 
with  which  we  are  acquainted.  It,  therefore,  seemed  best  to 
compare  the  atomic  weights  of  the  elements  directly  with  the 
atomic  weight  of  oxygen,  which  we  will  call  16,  since  it  is  ap- 
proximately 16  times  the  atomic  weight  of  hydrogen,  and  then 
compare  oxygen  with  hydrogen,  with  which  it  forms  the  very 
stable  compound,  water. 

Table  of  Atomic  Weights.  —  The  most  probable  atomic 
weights  of  the  elements,  based  upon  the  best  determinations, 
are  given  in  the  following  table :  — 


ELEMENT 

ATOMIC 
WEIGHT 

ELEMENT 

ATOMIC 
WEIGHT 

Aluminium     .... 

27.1 

Cadmium 

112  35 

Antimony            , 

120  0 

CsBsium 

132  9 

Argon    .... 

39.9 

Calcium 

40  1 

Arsenic  . 

75  0 

Carbon 

12  0 

Barium  .                       * 

137  4 

Cerium 

140  0 

Bismuth      

208.3 

Chlorine    

35.45 

Boron     

11.0 

Chromium      .... 

52.1 

Bromine     

79.96 

Cobalt  

59.0 

122 


ELEMENTS   OF   INORGANIC   CHEMISTRY 


ELEMENT 

ATOMIC 
WEIGHT 

ELEMENT 

ATOMIC 
WEIGHT 

Columbian!     .... 

94.0 

Platinum  .  '  . 

195  0 

Copper  . 

63.6 

Potassium  ... 

39  14 

Erbium  

166.0 

Praseodymium 

140  45 

Fluorine     .         ... 

19.05 

Rhodium                     .. 

103  0 

Gallium  . 

70.0 

Rubidium  . 

85  4 

Germanium    .... 
Glucinum 

72.5 
9  1 

Ruthenium    .... 
Samarium 

101.7 
150  0 

Gold  

197.25 

Scandium  

44  1 

Helium  

4.0 

Selenium  . 

79  2 

Hydrogen  . 

1  01 

Silicon  . 

28  4 

Indium  . 

114  0 

Silver    .      • 

107  93 

//Iodine 

126  85 

Sodium 

23  05 

Iridium  

193.0 

Strontium      .... 

87  68 

Iron  

65.9 

Sulphur     

32  06 

Krypton 

81  75 

Tantalum  . 

183  0 

Lanthanum 

138  8 

Tellurium  . 

127  6 

Lead. 

206  9 

Terbium 

160  0 

Lithium 

7  03 

Thallium 

204  1 

Magnesium     .... 
Manganese 

24.36 
55  0 

Thorium    
Thulium    

232.5 
171  0? 

Mercury 

200.0 

Tin   .     . 

119  0 

Molybdenum 

96  0 

Titanium 

48  15 

Neodymium 

143  6 

Tungsten 

184  0 

Neon      

20.0 

Uranium    

238.5 

Nickel    ...... 

58.7 

Vanadium     .... 

51.4 

Nitrogen 

14  04 

Xenon  . 

128.0 

Osmium 

191  0 

Ytterbium 

173  0 

Oxygen 

16  0 

Yttrium 

89  0 

Palladium  

106.5 

Zinc  

65.4 

Phosphorus     .... 

31.0 

Zirconium  

90.6 

The  Quantitative  Method  of  Dealing  with  Chemical  Reactions. 

—  Thus  far  we  have  dealt  qualitatively  with  chemical  re- 
actions— have  studied  what  substances  react  and  what  prod- 
ucts are  formed.  We  have  not  raised  the  question  as  to  how 
much  of  each  substance  entered  into  the  reaction  and  how 
much  of  each  product  was  formed. 


DETERMINATION   OF  RELATIVE  ATOMIC   WEIGHTS      123 

Knowing  the  relative  weights  of  the  atoms,  we  can  now  deal 
quantitatively  with  reactions,  and  from  given  quantities  of 
substances  calculate  the  amounts  of  the  products  formed.  We 
shall  now  apply  this  method  to  some  of  the  reactions  already 
studied. 

PROBLEMS  IN  CONNECTION  WITH  OXYGEN 

1.  Calculate  the  weight  of  oxygen  obtainable  from  4  grams 
of  potassium  chlorate.1 

2.  What  would  be  the  volume  of  oxygen  obtainable  from  5 
grams  of  manganese  dioxide  ?     (One  litre  of  oxygen  weighing 
1.4296  grams.) 

3.  How  many  grams  of  mercury  oxide  would  be  required  to 
yield  20  litres  of  oxygen  ?     How  many  to  yield  20  grams  of 
oxygen  ? 

4.  What  would  be  the  weight  and  volume  of  oxygen  required 
to  combine  with  10  grams  of  sulphur  to  form  sulphur  dioxide  ? 

5.  What  would  be  the  weight  of  phosphorus  required  to 
combine  with  15  litres  of   oxygen   to  form  phosphorus  pen- 
toxide  ?     What  would  be  the  weight  of  the  product  formed  ? 

6.  With  what  volume  of  oxygen  would  5  grams  of  carbon 
combine  to  form  carbon  dioxide  ? 


PROBLEMS  IN  CONNECTION  WITH  HYDROGEN 

1.  Weight  and  volume  of  hydrogen  that  can  be  obtained  by 
decomposing  5  grams  of  water  with  sodium?  (One  litre  of 
hydrogen  weighing  0.08995  gram.)  By  dissolving  1  gram  of 
zinc  in  hydrochloric  acid  ? 

1  The  solution  of  such  problems  is  comparatively  simple.  Potassium 
chlorate  has  the  formula  KC103.  The  molecular  weight  of  this  compound 
is  obtained  by  adding  the  atomic  weights  of  the  atoms  in  the  compound. 

At.  wt.  K  =  39.14 

At.  wt.  Cl=  35.45 

At.  wt.  0  =  16  x  3  =  48.00 

122.59 

The  molecular  weight  of  potassium  chlorate  is  to  the  atomic  weight  of 
oxygen  multiplied  by  the  number  of  atoms  of  oxygen  set  free  from  each 
molecule  of  the  salt  as  the  weight  of  potassium  chlorate  is  to  x :  — 

122,59;  48  =4  :«. 


124  ELEMENTS   OF  INORGANIC   CHEMISTRY 

2.  What  volume  of  hydrogen  will  combine  with  5  grams  of 
oxygen  to  form  water? 

3.  Weight    of   copper   oxide    which   would   be   reduced   to 
metallic  copper  by  3  litres  of  hydrogen  gas  ? 

4.  What  volume  of  hydrogen  would  combine  with  oxygen  to 
form  3  grams  of  water  ? 

5.  How  much  hydrogen  (weight  and  volume)  would  be  ob- 
tained by  electrolyzing  2  grams  of  water  ? 

6.  Volume  of  water-vapor  formed  when  3  grams  of  oxygen 
combine  with  hydrogen? 


PROBLEMS   IN   CONNECTION   WITH   CHLORINE 

1.  Weight  and  volume  of  chlorine  obtainable  from  1  gram 
of  hydrochloric  acid  ?     (One  litre  of  chlorine  weighing  3.22 
grams.) 

2.  Weight  of  copper  which  would  combine  with  1  litre  of 
chlorine  ? 

3.  Volume  of  chlorine  which  would  combine  with  1  litre  of 
hydrogen  ?  with  1  gram  of  hydrogen  ?     What  would  be  the 
volume  of  the  hydrochloric  acid  formed  in  each  case  ? 

4.  How  much  sodium  chloride  would  be  required  to  yield 
the  chlorine  which  would  combine  with  2  grams  of  hydrogen  ? 

5.  Weight  of  potassium  chlorate  which  could  be  formed  from 
all  the  chlorine  in  3  grains  of  bleaching-powder  ? 


PROBLEMS  IN  CONNECTION  WITH  NITROGEN 

1.  When  3  grams  of  ammonium  nitrite  are  decomposed  by 
heat,  what  is  the  weight  of  the  nitrogen  obtained  ?     What  is 
the  volume  ?     (One  litre  of  nitrogen  weighing  1.2575  grams.) 

2.  With  what  volume  of  ammonia  would  2  grams  of  hydro- 
chloric acid  combine  ?     (One  litre  of  ammonia  weighing  0.775 
grams.) 

3.  What  amount  of  ammonium  nitrite  must  be  decomposed 
by  heat  to  yield  one  gram  of  nitrous  oxide  ? 

4.  When  5  grams  of  sodium  nitrite  are  decomposed  by  sul- 
phuric acid,  what  weight  of  nitrogen  sesquioxide  is  set  free? 

5.  What  would  be  the  weight  of  the  nitric  acid  formed  by 
treating  7  grains  of  sodium  nitrate  with  the  quantity  of  sul- 
phuric acid  required  to  produce  the  acid  sulphate,  NaHS04? 


DETERMINATION  OF   RELATIVE   ATOMIC    WEIGHTS      125 

PROBLEMS   IN   CONNECTION   WITH  THE  VOLUME 
OF   GASES 

1.  If  the  volume  of  a  dry  gas  at  0°  and  729  mm.  pressure  is 
50  cc.,  what  would  be  its  volume  under  normal  conditions  (i.e. 
0°  and  760  mm.  pressure)  ?  1 

2.  If  the  volume  o£  a  dry  gas  at  10°  and  760  mm.  pressure  is 
25  cc.,  what  would  be  its  volume  under  normal  conditions  ? 

3.  If  the  volume  of  a  gas  over  water  at  20°  and  745  mm. 
pressure  is  35  cc.,  what  volume  would  it  have  at  normal  tem- 
perature and  pressure  (the  pressure  of  the  aqueous  vapor  at  20° 
being  17.4  mm.)  ? 

1  The  solution  is  as  follows  :  — 

The  equation  for  the  volume  of  a  gas  under  normal  conditions,  as  de- 

veloped  on  page  112,  is     r= 


v  =  50, 

p  =  729, 

h  =  0,    since  the  gas  is  dry  ; 
«  =  0. 
Substituting  these  values  in  the  above  equation,  we  have  :  — 

v_     50(729) 
"  760(1  +  0)' 


CHAPTER  XIII 

THE  PERIODIC   SYSTEM 

Relations  between  Atomic  Weights  and  Properties  of  the 
Elements.  —  Almost  as  soon  as  the  atomic  weights  of  the  ele- 
ments were  determined,  chemists  began  to  look  for  relations 
between  the  atomic  weights  of  the  elements  and  their  proper- 
ties. A  number  of  such  relations  were  pointed  out,  however 
of  not  very  wide-reaching  significance. 

The  Periodic  System.  —  The  first  to  point  out  the  most 
important  features  in  the  arrangement  of  the  elements  accord- 
ing to  their  atomic  weights  was  the  Russian,  Mendeleeff.  In 
1869  he  arranged  the  elements  in  a  table  in  the  order  of  their 
atomic  weights,  and  showed  clearly  that  there  is  a  periodic 
recurrence  of  properties  as  the  atomic  weights  increase.  This 
will  be  seen  best  in  the  following  table,  which  is  essentially  the 
same  in  principle  as  that  of  Mendeleeff.  Some  elements 
unknown  at  the  time  the  original  system  was  proposed  are 
introduced,  and  the  present  atomic  weights  of  the  elements  are 
given. 

The  elements  in  group  0  are  rare  elements  occurring  in  the 
air,  and,  as  we  have  seen,  have  been  recently  discovered.  The 
elements  scandium,  gallium,  and  germanium  were  unknown 
when  Mendeleeff  proposed  his  table.  He,  however,  predicted 
their  existence,  their  properties,  and  the  properties  of  their 
compounds.  These  predictions  have  been  fulfilled  in  a  most 
beautiful  manner,  and  this  is  one  of  the  crowning  glories  of 
the  periodic  system. 

All  the  elements  are  arranged  in  succession  in  the  order  of 
their  increasing  atomic  weights.  If  we  start  with  the  element 

126 


t-H  <1 

o  o" 

7  7 

O  £ 


a.  W  0 


1     1 


§  M  M 
O 


8 
II        fe 


T-<  fr- 


S    3 


12. 


127 


128  ELEMENTS  OF   INORGANIC   CHEMISTRY 

with  the  smallest  atomic  weight  next  to  hydrogen,  i.e.  lithium, 
and  arrange  the  succeeding  elements  in  the  order  of  their  atomic 
weights  up  to  fluorine,  we  find  that  the  next  element,  sodium, 
has  properties  quite  similar  to  those  of  lithium.  If  we  place 
sodium  in  the  same  vertical  column  with  lithium,  and  then 
arrange  the  next  elements  in  the  order  of  their  atomic  weights, 
we  find  that  magnesium  falls  in  the  same  column  with  glucinum 
or  beryllium,  aluminium  with  boron,  silicon  with  carbon,  phos- 
phorus with  nitrogen,  sulphur  with  oxygen,  and  chlorine  with 
fluorine.  This  is,  of  course,  a  remarkable  relation,  since  in 
every  case  those  elements  that  fall  in  the  same  vertical  column 
resemble  each  other  very  closely.  The  first  seven  elements, 
starting  (not  with  hydrogen,  since  it  does  not  fit  into  this 
scheme)  with  lithium,  and  ending  with  fluorine,  agree  very 
closely  in  properties  with  the  second  set  of  seven  elements 
arranged  as  in  the  above  table.  We  come  now  to  the  first 
member  of  the  next  series  of  seven  elements,  —  potassium ;  it 
falls  right  into  the  group  with  lithium  and  sodium,  calcium  with 
glucinum  and  magnesium,  titanium  with  carbon  and  silicon, 
vanadium  with  nitrogen  and  phosphorus,  chromium  with  oxygen 
and  sulphur,  and  manganese  with  fluorine  and  chlorine.  Here 
again  striking  analogies  appear  between  the  different  members 
in  the  same  groups.  The  blank  space  between  calcium  and 
titanium  contained  no  known  element  when  this  table  was 
prepared.  The  element  has  since  been  discovered.  After  we 
leave  manganese  we  encounter  one  of  the  weakest  points  of 
the  Periodic  System.  The  next  elements  in  order  of  atomic 
weights  are  iron,  cobalt,  and  nickel;  but  it  is  obvious  that 
neither  of  these  can  be  placed  in  the  same  group  with  the 
alkali  metals.  They  must,  therefore,  be  set  aside  arid  left  out 
of  the  system.  Then  we  come  to  copper,  which  is  very  ques- 
tionably placed  with  the  members  of  group  I.  Then  irregu- 
larities appear  again.  At  the  end  of  the  sixth  series  we  find 
three  or  four  more  elements  which  do  not  fit  into  the  scheme, 
but  after  leaving  these,  regularities  again  begin  to  manifest 
themselves. 


THE   PERIODIC   SYSTEM  129 

The  above  suffices  to  show  the  general  relation,  and  also  the 
periodic  recurrence  of  properties  with  increase  in  the  atomic 

weights. 

CHEMICAL   PROPERTIES  AND  ATOMIC   WEIGHTS. 
COMBINING   POWER 

If  we  start  with  lithium  and  proceed  to  the  right  along  the 
second  series,  this  striking  fact  is  observed:  the  elements  in- 
crease in  their  power  to  combine  with  oxygen  regularly  from 
left  to  right.  Take  first  the  power  of  the  elements  to  combine 
with  oxygen.  Lithium  forms  the  compound  Li20,  beryllium 
BeO,  aluminium  A1203,  carbon  C02,  nitrogen  N205;  oxygen  and 
fluorine  may  be  disregarded  for  the  moment.  Take  the  third 
series.  Sodium  forms  the  compound  NaO  which  is  a  super- 
oxide,  magnesium  MgO,  aluminium  A1203,  silicon  Si02,  phos- 
phorus P205,  sulphur  S03,  and  chlorine  C1207.  The  fourth  and 
fifth  series  show  the  same  regularities,  and  similar  relations  are 
observed  throughout  the  table.  The  best  example  of  an  element 
octivalent  towards  oxygen  is  osmium,  which  forms  the  com- 
pound Os04.  We  have,  then,  NaO,  MgO,  A1203,  Si02,  P205,  S03, 
C1207,  Os04. 

We  may  say  in  general  that  the  power  of  the  elements  to 
combine  with  oxygen  is  smallest  in  group  I,  and  increases 
regularly  by  unity  in  each  succeeding  group ;  reaching  a  maxi- 
mum in  group  VIII,  where,  at  least  in  the  case  of  osmium,  it  is 
eight. 

Results  of  a  similar  character  are  obtained  if  we  study  the 
power  of  the  elements  to  combine  with  chlorine.  Sodium  com- 
bines with  one  chlorine  atom,  magnesium  with  two,  aluminium 
with  three,  silicon  with  four,  phosphorus  with  five.  Sulphur 
does  not  combine  directly  with  six  chlorine  atoms,  but  combines 
with  both  oxygen  and  chlorine,  forming  the  compound  S02C12, 
in  which  the  sulphur  has  a  valence  of  four  towards  the  oxygen, 
and  of  two  towards  the  chlorine,  or  of  six  in  all.  But  there  is  a 
member  of  group  VI  which  combines  directly  with  six  chlorine 


130  ELEMENTS   OF  INORGANIC   CHEMISTRY 

*    •»  * 

atoms.  This  is  tungsten,  in  the  tenth  series.  We  would 
express  the  combining  power  of  the  elements  towards  chlorine 
as  follows  :  — 

NaCl,    MgCl2,    A1C13,    SiCl4,    PC15,    S02C12. 

(WC16) 

Exactly  the  same  regularity  which  was  observed  in  the  case  of 
oxygen  exists  here.  The  elements  in  group  I  have  the  smallest 
power  to  combine  with  chlorine,  and  this  increases  by  unity 
from  group  to  group  as  we  pass  from  left  to  right  ;  reaching  a 
maximum  of  six  in  the  sixth  group.  We  know  of  no  element 
which  has  the  power  of  combining  directly  with  more  than  six 
atoms  of  chlorine. 

When  we  examine  the  power  of  the  elements  to  combine  with 
hydrogen,  a  regularity  is  observed,  but  of  a  different  kind  from 
those  already  considered.  Some  of  the  elements  in  groups  I, 
II,  and  III  combine  with  hydrogen,  forming  hydrides,  and 
these  have  the  compositions  in  general  indicated  in  the  table. 
When  we  come  to  group  IV,  we  find  in  carbon  a  remark- 
able power  to  combine  with  hydrogen.  The  highest  valence  of 
the  elements  towards  hydrogen  is  manifested  in  this  group, 
where  one  atom  of  the  element  combines  directly  with  four 
atoms  of  hydrogen.  As  we  pass  to  the  right,  the  power  of  the 
elements  to  combine  with  hydrogen  decreases,  and  decreases 
regularly.  Nitrogen  combines  with  three  atoms  of  hydrogen, 
oxygen  with  two,  and  fluorine  with  one  :  — 


NaH,    CaH2,    BH3,    CH4,    NH3,    OH2,    FH. 

The  valence  towards  hydrogen  increases  from  group  I  to 
group  IV,  where  it  reaches  a  maximum,  and  diminishes  regu- 
larly as  the  valence  towards  oxygen  increases. 

The  relations  pointed  out  between  the  combining  power  of  the 
elements  are  general,  extending  throughout  the  table  of  the 
elements.  It  should,  however,  be  stated  here  that  there  are 
many  breaks  in  the  system,  irregularities  appearing  in  many 
places. 


THE  PERIODIC   SYSTEM  131 

<d^^ 

Basic  ™d  Add  Properties 

At  least  one  other  relation  between  the  chemical  properties 
of  the  elements  and  their  atomic  weights  must  be  pointed  out. 
In  any  given  series  the  element  with  the  lowest  atomic  weight 
has  thjejmiallest  J>ower  to_combine _with_oxygen,  as  has  already 
'  been  stated.  It  Has  also  the  strongest  basic  character.  Thus, 
lithium  is  more  basic  than  glucinum,  which,  in  turn,  is  far 
more  basic  than  boron.  Sodium  is  more  basic  than  magnesium, 
while  aluminium  begins  to  show  acid  properties  in  its  hy- 
droxide. Potassium  is  far  more  basic  than  calcium,  rubidium 
than  strontium,  caesium  than  barium.  The  difference  between 
copper  and  zinc,  and  silver  and  cadmium,  is  not  so  striking. 
As  we  find  the  most  basic  elements  in  the  first  group,  we  would 
expect  to  find  the  most  acid  in  the  last,  and  such  is  the  case. 
Through  the  middle  groups  we  find  elements  which  show,  now 
more,  now  less,  basic  or  acid  properties,  depending  upon  condi- 
tions ;  but  in  the  last  column  of  the  last  well-defined  group  we 
have  elements  which  manifest  only  acid-forming  properties. 
The  hydrogen  and  hydroxyl  compounds  of  the  halogens  are 
always  acids,  and  always  react  as  such  with  all  other  sub- 
stances. These  facts  are  very  surprising.  As  we  pass  upward 
in  the  table  of  atomic  weights,  say  from  oxygen,  the  first  ele- 
ment we  encounter  is  fluorine,  with  very  pronounced  acid- 
forming  properties.  The  element  with  the  next  higher  atomic 
weight  is  sodium,  which  is  one  of  the  strongest  base-forming 
elements.  Similarly,  next  to  sulphur  comes  chlorine,  which 
has  stronger  acid-forming  properties  than  sulphur,  but  next 
to  chlorine  comes  potassium,  which  is  one  of  the  most  strongly 
basic  elements.  In  the  same  way  bromine  is  followed  by  rubid- 
ium, and  iodine  by  caesium,  where  the  contrast  in  properties 
is  quite  as  great  as  in  the  cases  referred  to  above. 

Many  other  relations  between  chemical  properties  and  atomic 
weights  have  been  pointed  out,  but  those  already  considered 
are  among  the  most  important. 


132  ELEMENTS   OF  INORGANIC   CHEMISTRY 

General  Scheme  to  be  Followed.  —  In  dealing  with  the  re- 
maining elements  we  shall  be  guided  largely  by  the  Periodic 
System.  This  system,  however,  is  defective,  and  we  shall, 
therefore,  not  follow  it  blindly,  but  depart  from  it  whenever 
the  relations  can  be  more  clearly  established  by  doing  so. 

We  shall  begin  with  the  members  of  group  VII  in  Men- 
deleeff's  table,  omitting,  however,  one  member,  manganese, 
which  will  not  be  taken  up  until  much  later. 

We  shall  then  take  up  some  of  the  members  of  group  VI  — 
sulphur,  selenium,  and  tellurium ;  while  chromium,  molybde- 
num, tungsten,  and  uranium  will  not  be  referred  to  until  very 
much  later.  The  nitrogen  group  (V)  will  then  be  studied,  and 
following  this  the  carbon  group  (IV). 

The  metallic  elements  will  then  be  taken  up.  Groups  I  and 
II  will  be  studied  very  nearly  in  the  order  indicated  in  the 
Periodic  System,  while  the  remaining  metallic  elements  will  be 
studied  more  or  less  independent  of  the  system. 


CHAPTER  XIV 

BROMINE,  IODINE,  FLUORINE 

BROMINE  (At.  Wt.  =  79.96) 

Occurrence  and  Preparation.  —  The  element  bromine  closely 
resembles  the  element  chlorine.  Like  the  latter  it  does  not 
occur  in  the  free  condition  in  nature,  and  exists  in  very  much 
smaller  quantity  than  chlorine.  The  compounds  of  bromine, 
like  those  of  chlorine,  being  in  general  very  soluble,  the  chief 
occurrence  of  bromine  is  in  the  waters  of  the  sea.  Where  sea- 
water  has  evaporated  and  deposited  the  great  salt  beds  of  the 
earth,  we  find  the  bromides  mixed  with  a  large  number  of  other 
salts. 

Bromine  is  prepared  by  three  methods :  The  electrolysis  of 
bromides,  which  is  strictly  analogous  to  the  electrolysis  of 
chlorides,  the  bromine  ion  passing  to  the  anode  and  separating 
in  the  free  condition,  while  the  metal  passes  to  the  cathode. 

A  second  method  which  was  much  used  formerly,  but  is  now 
seldom  employed  on  the  large  scale,  consists  in  the  oxidation 
of  hydrobromic  acid  by  manganese  dioxide ;  the  hydrobromic 
acid  being  set  free  from  the  bromide  by  means  of  sulphuric 
acid :  — 

,       2  NaBr  +  H2S04  =  Na2S04  +  2  HBr ; 
2  HBr  +  Mn02  +  H2S04  =  MnS04  +  2  H20  +  Br2. 

Combining  these  in  one  equation,  we  have  — 
2  NaBr  +  Mn02  +  2  H2S04  =  MnS04  +  Na2S04+  2  H20  +  Br2. 

The  third  method  consists  in  the  replacement  of  bromine 
from  bromides  by  means  of  chlorine  :  — 

2  K  Br  +  C12  =  2  KC1  +  2  Br. 

This  method  finds  extensive  use  to-day. 

133 


134  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Chemical  Properties  of  Bromine.  —  Bromine  in  its  chemica 
properties  strikingly  resembles  chlorine.  Like  the  latter  i 
unites  directly  with  most  of  the  elements.  The  compound 
formed  —  the  bromides  —  are  not  as  stable  as  the  chlorides 
This  is  shown  by  the  fact  that  chlorine  replaces  bromine  fron 
the  bromides.  The  bromides,  like  the  chlorides,  are  in  genera 
soluble  in  water,  the  bromides  being  more  soluble  than  th 
chlorides. 

Bromine  has  a  remarkable  power  to  disintegrate  organi 
substances.  Its  action  upon  the  mucous  membrane  of  th 
throat  and  nose  is  more  vigorous  than  that  even  of  chlorine 
and  great  precaution  must,  therefore,  be  taken  in  working 
with  the  substance,  to  be  protected  from  its  disintegratin 
fumes. 

Detection  of  Bromine.  —  The  chemical  properties  of  bromin 
are  so  closely  allied  to  those  of  chlorine  that  it  might  at  firs 
sight  seem  difficult  to  determine  with  which  we  are  dealing 
This  difficulty  is,  however,  only  apparent.  If  a  bromide  i 
treated  with  chlorine,  the  bromine,  as  we  have  seen,  is  se 
free  and  can  be  recognized  by  its  odor  and  color.  If  a  solutio: 
of  a  bromide  is  treated  with  a  little  chlorine  water,  a  littl 
carbon  bisulphide  being  added  to  the  tube  and  shaken  vigoi 
ously,  the  bromine  which  has  been  set  free  by  the  cnlorin 
is  dissolved  by  the  carbon  bisulphide,  and  imparts  its  charactei 
istic  reddish  brown  color  to  the  solution. 

Bromine  Atoms  and  Bromine  Ions.  —  Bromine  in  the  atomi 
condition  behaves  very  differently  from  bromine  in  the  ioni 
condition.  The  bromine  ions  when  brought  in  contact  wit; 
the  silver  ions  combine  with  them  at  once,  forming  the  charac 
teristic  precipitate,  silver  bromide. 

K,  Br  +  Ag,  NO,  =K,  NO,  +  AgBr. 

When  bromine  in  the  un-ionized  condition  is  brought  into  th 
presence  of  silver  ions,  it  does  not  combine  to  the  slightes 
extent  with  the  silver  ions. 


BROMINE,    IODINE,    FLUORINE  135 

Physical  Properties  of  Bromine.  —  The  reddish  brown  liquid, 
bromine,  has  a  density  3.1.  It  boils  at  63°,  yielding  a  reddish 
yellow  vapor.  The  density  of  the  vapor  is  160,  in  terms  of 
hydrogen  as  two.  The  atomic  weight  being  80,  the  molecule 
of  bromine  contains  two  atoms,  or  has  the  composition  repre- 
sented by  the  formula  Br2.  Bromine  solidifies  at  —  7°,  forming 
an  orange-red  solid. 

Bromine  dissolves  in  water,  forming  what  is  known  as 
bromine  water,  which  is  analogous  to  chlorine  water.  The 
saturated  solution  in  water  contains  about  3  per  cent  of 
bromine. 

Hydrobromic  Acid.  —  A  method  of  preparing  hydrobromic 
acid  is  by  allowing  water  to  act  upon  the  bromides  of  certain 
acid-forming  elements  such  as  phosphorus.  The  reaction  that 
takes  place  in  this  case  is  :  — 

PBr5  +  H20  =  POBr3  +  2  HBr, 

resulting  in  the   formation   of   phosphorus   oxybromide  and 
hydrobromic  acid. 

The  question  which  would  naturally  be  asked  is,  Why  not 
prepare  hydrobromic  acid  by  the  action  of  sulphuric  acid  on 
bromides  ?  This  would  be  analogous  to  the  preparation  of 
hydrochloric  acid  by  the  action  of  sulphuric  acid  on  chlorides. 

This  method  is  possible,  but  precautions  must  be  taken,  or  a 
secondary  reaction  between  the  hydrobromic  acid  formed  and 
the  sulphuric  acid  takes  place,  which  interferes  with  the  value 
of  the  method.  The  sulphuric  acid  may  be  reduced  to  sulphur 
dioxide :  — 

2  KBr  +  H2S04  =  K2S04  +  2  HBr; 
H2S04  +  2  HBr  =  2  H20  +  2  Br  +  S02. 

Properties  of  Hydrobromic  Acid.  —  Hydrobromic  acid  re- 
sembles hydrochloric  acid  very  closely  in  its  chemical  and 
physical  properties. 

Hydrobromic  acid  is  a  reducing  agent.  It  is  also  a  very 
strong  acid,  but  not  quite  as  strong  as  hydrochloric  acid,  A 


136  ELEMENTS   OF  INORGANIC   CHEMISTRY 

dilute,  aqueous  solution  of  hydrobromic  acid  is  completely  dis- 
sociated into  its  ions  :  — 

HBr  =  H,  Br. 

The  bromine  ions  manifest  their  presence  by  combining  with 
the  silver  ions  when  brought  in  contact  with  them,  forming  the 
white  precipitate,  silver  bromide,  which  is  soluble  with  diffi- 
culty in  ammonia. 

When  chlorine  is  brought  into  the  presence  of  a  bromide, 
the  bromine  separates,  and  chlorine  passes  into  solution.  If  we 
examine  this  reaction  more  closely,  we  find  that  what  has  taken 
place  is  a  transfer  of  the  electrical  charge  from  the  bromine 
ion  to  the  chlorine  atom,  converting  the  latter  into  an  ion, 
while  the  bromine  ion  having  lost  its  charge  is  converted  into 
an  atom.  The  reaction  that  takes  place  would  be  represented 
thus  :  — 


H,  Br  +  Cl  =  H,  Cl  +  Br. 

Compounds  of  Bromine  with  Oxygen  and  Hydrogen.  —  Bromine 
forms  two  well-characterized  acids  with  oxygen  and  hydrogen. 
These  are  hypobromous  and  bromic  adds.  While  hydrobromic 
acid  is  less  stable  than  hydrochloric,  as  we  have  seen,  oxygen 
acids  of  bromine  are  more  stable  than  the  corresponding  acids 
of  chlorine. 

The  sodium  salt  of  hypobromous  acid  is  prepared  by  the 
action  of  bromine  on  sodium  hydroxide  :  — 

2  NaOH  +  2  Br  =  H20  +  NaBr  +  NaOBr. 

This  method  of  preparing  sodium  hypobromite  is  strictly 
analogous  to  the  method  of  preparing  sodium  hypochlorite. 

The  best  method  of  preparing  bromic  acid,  is  by  the  action 
of  bromine  on  caustic  potash.  A  mixture  of  potassium  bro- 
mide and  bromate  is  formed  :  — 

6  KOH  +  3  Br2  =  5  KBr  +  KBr03  +  3  H20. 

From  this  mixture  the  potassium  bromate  can  be  readily 
separated  by  fractional  crystallization,  —  the  same  method 


BROMINE,   IODINE,    FLUORINE  137 

which  was  employed  to  separate  potassium  chlorate  from 
potassium  chloride.  The  brornate  is  much  less  soluble  in 
water  than  the  bromide,  and  readily  crystallizes  from  a  not  too 
dilute  solution  of  the  two  salts  in  water. 

Bromic  acid  is  obtained  from  potassium  bromate  by  methods 
strictly  analogous  to  those  employed  for  obtaining  chloric  acid 
from  potassium  chlorate. 

IODINE  (At.  Wt.  =  126.85) 

Occurrence  and  Preparation.  —  Iodine,  named  from  the  violet- 
blue  color  of  its  vapor,  occurs  very  rarely  in  the  free  condition. 
It  is  widely  distributed  in  nature,  existing,  however,  usually 
only  in  small  quantities.  It  occurs  along  with  chlorides  and 
bromides,  but  in  very  much  smaller  quantities  than  either  of 
these  substances.  It  is  also  contained  in  sea-water,  as  we  would 
expect,  on  account  of  the  solubility  of  its  compounds.  It  is  taken 
up  from  the  waters  of  the  sea  by  certain  sponges  and  plants, 
and  exists  in  considerable  quantity  in  the  ashes  of  such  plants. 

Iodine  is  obtained  to-day  mainly  from  Chili  saltpetre,  in 
which  it  occurs  in  the  form  of  sodium  iodate.  The  iodine  is 
obtained  from  this  salt  by  reduction  with  sulphurous  acid. 

When  an  iodide  is  treated  with  manganese  dioxide  and  sul- 
phuric acid,  the  following  reaction  takes  place  :  — 

2  KI  +  Mn02  +  2  H2S04  =  MnS04  +  K2S04  +  2  H20  +  !«. 

Iodine  can  also  be  displaced  from  iodides  by  means  of  chlorine. 
When  a  solution  of  potassium  iodide  is  treated  with  chlorine 
water,  the  following  reaction  takes  place :  — 

K,  I  +  Cl  =  K,  Cl  +  I. 

This  is  analogous  to  the  action  of  chlorine  upon  bromides, 
which  we  have  seen  consists  in  a  transference  of  the  electrical 
charge  from  the  bromine  ion  to  the  chlorine.  Here  we  have 
the  charge  transferred  from  the  iodine  ion  to  the  chlorine, 
which  becomes  an  ion,  the  iodine  having  lost  its  charge 
becoming  an  atom. 


138  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Chemical  Properties  of  Iodine.  —  The  purplish  black  solid, 
iodine,  resembles  strongly  in  its  chemical  properties  the  ele- 
ments chlorine  and  bromine.  It  combines  with  many  other 
elements,  but  is  not  quite  as  active  chemically  as  bromine  and 
chlorine.  Iodine  forms  a  characteristic  blue  color  with  starch 
paste,  which  enables  its  presence  to  be  easily  detected. 

If  the  iodine  is  combined  as  in  an  iodide,  it  can  be  detected 
by  simply  adding  concentrated  sulphuric  acid.  The  hydriodic 
acid  set  free  reacts  with  the  sulphuric  acid,  reducing  it,  and 
the  iodine  is  liberated  as  such.  It  can  be  recognized  by  the 
characteristic  dark  brown  color  which  it  imparts  to  the  solution. 

Physical  Properties  of  Iodine.  —  Although  iodine  is  a  solid 
at  ordinary  temperatures,  it  can  readily  be  converted  into  vapor 
at  a  slightly  elevated  temperature.  When  iodine  is  heated 
exposed  to  the  air,  i.e.  under  ordinary  conditions,  it  does  not 
melt,  but  passes  at  once  into  vapor.  It  can,  however,  be 
melted  at  114°,  and  boils  at  about  185°. 

.  The  vapor  of  iodine  can  be  readily  condensed  to  a  solid 
when  the  temperature  of  the  vapor  is  again  lowered.  This 
process  of  converting  a  solid  into  a  vapor,  and  recondensing 
the  vapor  to  a  solid,  is  known  as  sublimation. 

Iodine  dissolves  in  water  to  only  a  slight  extent.  If  to  the 
water  potassium  iodide  or  hydriodic  acid  is  added,  the  solution 
dissolves  iodine  in  considerable  quantities.  Iodine  dissolves 
readily  in  carbon  bisulphide,  chloroform,  alcohol,  and  ether. 
Solutions  of  iodine  in  the  last  two  solvents  are  known  as 
tincture  of  iodine. 

Hydriodic  Acid,  HI. — When  a  mixture  of  hydrogen  and 
iodine  is  heated,  there  is  partial  combination  between  the  two, 
forming  hydriodic  acid,  HI.  Only  a  portion  of  the  mixture, 
however,  combines.  The  velocity  of  the  reaction,  i.e.  the  amount 
of  hydriodicr  acid  which  is  formed  in  a  given  time,  can  be 
greatly  increased  by  heating  the  mixture  in  the  presence  of 
finely  divided  platinum,  which  acts  by  contact,  or  catalyti catty, 
as  we  say.  Even  under  these  conditions  a  complete  combina- 
tion of  the  two  gases  cannot  be  effected. 


BROMINE,   IODINE,   FLUORINE  139 

A  far  more  convenient  method  of  preparing  hydriodic  acid 
is  by  the  action  of  water  on  phosphorus  triiodide,  PI3:  — 

PI3  +  3  H20  =  H3P03  +  3  HI. 

Hydriodic  acid  cannot  be  prepared  by  the  action  of  sulphuric 
acid  on  iodides,  since,  as  we  have  seen,  hydriodic  acid  reduces 
sulphuric  acid.  The  extent  to  which  this  reduction  takes 
place  depends  largely  upon  the  temperature  and  concentration 
of  the  solutions,  hydrogen  sulphide  being  usually  formed :  — 

2  KI  +  H2S04  =  K2S04  +  2  HI ; 

H2S04+  8  HI  =  4  H20  +  81  +  H2S. 

Compounds  of  Iodine  with  Oxygen  and  Hydrogen.  —  Iodine 
forms  two  well-characterized  compounds  with  oxygen  and 
hydrogen.  These  are  iodic  acid  and  periodic  acid.  When 
iodine  is  dissolved  in  caustic  potash,  the  reaction  takes  place 
thus :  — 

6  KOH  +  6  I  =  5  KI  +  KI03  +  3  H20. 

The  iodates  are  more  readily  prepared  by  the  action  of 
iodine  on  chlorates.  When  no  water  is  present,  we  have :  — 

2  KC103  +  21  =  2  KI03  +  C12. 

Periodic  acid  does  not  have  a  composition  strictly  analogous 
to  perchloric  or  perbromic  acid,  but  this  plus  two  molecules  of 
water :  — 

HI04  +  2  H20  =  H5I06. 

(unknown) 

FLUORINE  (At.  Wt.  =  19.05) 

Occurrence  and  Preparation.  —  Fluorine,  on  account  of  its 
unusual  chemical  activity,  does  not  occur  in  nature  in  .the  free 
condition.  It  occurs  mainly  in  combination  with  the  element 
calcium  as  fluor-spar,  from  which  it  derives  its  name.  Fluor- 
spar is  so  called  because  it  readily  melts  and  flows,  serving  as 
a  flux  for  other  substances.  Fluorine  also  occurs  in  another 
mineral,  cryolite,  in  considerable  quantity.  -  Cryolite  is  a 


140 


ELEMENTS   OF   INORGANIC   CHEMISTRY 


double  fluoride  of  sodium  and  aluminium,  occurring  mainly  in 
Greenland,  and  having  the  composition  Na3AlF6. 

The  problem  of  isolating  fluorine  remained  for  a  long  time 
unsolved.  On  account  of  the  great  chemical  activity  of  this 
substance,  as  soon  as  it  was  set  free  from  its  compounds  it 


FIG.  36. 

would  combine  again  with  whatever  it  came  in  contact.  The 
problem  of  isolating  fluorine  was  solved  by  the  French  chemist, 
Moissan. 

He  found  that  when  potassium  fluoride,  is  dissolved  in  an- 
hydrous hydrofluoric  acid  the  solution  conducts  the  current, 
hydrogen  being  liberated  at  the  cathode  and  fluorine  at  the 
anode.  At  first  the  attempt  was  made  to  use  vessels  lined: 


BROMINE,   IODINE,   FLUORINE  141 

with  calcium  fluoride  where  the  fluorine  escapes,  but  vessels  of 
platinum  were  subsequently  employed  and  found  to  work  very 
satisfactorily.  Indeed,  it  has  subsequently  been  shown  that 
fluorine  does  not  act  very  vigorously  upon  copper,  and  copper 
vessels  have  been  used  in  which  to  liberate  the  element 
fluorine. 

The  apparatus  used  by  Moissan  for  preparing  fluorine  is 
shown  in  Figure  36.  The  apparatus  and  electrodes,  made  of 
platinum-indium,  have  the  form  shown  in  the  figure.  The 
electrodes  are  insulated  by  means  of  stoppers  (S.S.)  of  fluor- 
spar. The  apparatus  is  kept  at  a  temperature  of  —  23°  by 
means  of  methyl  chloride.  The  fluorine,  liberated  on  the 
anode,  passes  out  through  a  spiral  platinum  tube  cooled  to 
—  50°  by  means  of  methyl  chloride,  and  then  through  two 
tubes  of  sodium  fluoride  to  remove  all  traces  of  hydrofluoric 
acid. 

Chemical  Properties  of  Fluorine.  —  Fluorine  is  one  of  the 
most  active  chemically  of  all  the  elements.  It  replaces  chlo- 
rine from  chlorides  and  from  hydrochloric  acid.  At  ordinary 
temperatures  it  combines  with  most  of  the  metals,  converting 
them  into  fluorides.  Platinum  and  gold  are  about  the  only 
metals  which  resist  its  action,  and  these  are  transformed  into 
fluorides  at  elevated  temperatures.  About  the  only  elements 
which  resist  the  action  of  fluorine  are  nitrogen,  chlorine, 
oxygen,  and  argon  and  its  associates. 

Physical  Properties  of  Fluorine.  —  Fluorine  is  a  gas  at  all 
ordinary  temperatures,  having  a  light,  greenish  yellow  color. 
It  is  much  lighter  in  color  than  chlorine,  and  much  more  active 
upon  the  mucous  membrane  of  the  nose  and  throat. 

Fluorine  has  been  liquefied  by  the  combined  efforts  of  Moissan 
in  France  and  Dewar  in  England.  The  fluorine  was  cooled  to 
— 190°  by  means  of  liquid  air,  when  it  liquefied,  and  was 
received  in  a  glass  bulb  with  a  vacuum-jacket.  Fluorine  boils 
at  — 187°,  and  at  this  low  temperature  has  lost  much  of  its 
chemical  activity,  as  is  obvious  from  the  fact  that  it  can  be 
received  in  a  glass  vessel.  Fluorine  has  recently  been  solidified. 


142  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Hydrofluoric  Acid,  HF.  —  Hydrogen  fluoride,  or  hydrofluoric 
acid,  is  prepared  most  conveniently  by  the  action  of  sulphuric 
acid  on  calcium  fluoride  :  — 

CaF2  +  H2S04  =  CaS04  +  2  HF. 

The  most  characteristic  chemical  property  of  hydrofluoric  acid 
is  its  power  to  act  upon  glass,  etching  it,  as  we  say.  It  is  ex- 
tensively used  for  this  purpose  in  preparing  measuring  appa- 
ratus especially  for  chemical  work.  Hydrofluoric  acid  does 
not  act  upon  paraffine  and  similar  organic  substances.  The 
glass  vessel  upon  which  it  is  desired  to  make  a  permanent  line 
is  covered  with  paraffine  by  dipping  it  into  the  molten  material. 
A  fine  line  is  then  drawn  through  the  paraffine  at  the  place 
where  it  is  to  appear  on  the  glass.  The  glass  is  thus  exposed 
at  this  place.  It  is  now  subjected  to  the  action  of  the  fumes 
of  hydrofluoric  acid.  Where  the  glass  is  protected  by  the  paraf- 
fine, it  is  not  acted  upon  by  the  fumes  of  the  acid.  Where  the 
paraffine  has  been  removed,  however,  the  glass  is  etched. 

Comparison  of  the  Several  Acids  formed  by  the  Halogens. — 
We  have  seen  that  chlorine,  bromine,  iodine,  and  fluorine  all 
form  compounds  with  hydrogen  which  are  acid.  Taking  these 
in  the  order  of  the  increasing  atomic  weight  of  the  halogen,  we 
have  seen  that  hydrofluoric  acid  is  the  most  stable  of  all  the 
compounds  of  the  halogens  with  hydrogen.  Hydrochloric  acid 
is  next,  and  this  is  followed  in  the  order  of  decreasing  stability 
by  hydrobromic  and  hydriodic  acids.  Indeed,  the  last-named 
substance  is  quite  unstable.  Of  the  four  acids ;  hydrochloric, 
hydrobromic,  and  hydriodic  are  of  about  equal  strength,  and 
are  among  the  strongest  acids,  while  hydrofluoric  acid  is  a 
much  weaker  acid. 

If  we  turn  to  the  compounds  with  oxygen  and  hydrogen,  we 
find  the  order  exactly  reversed.  Fluorine  forms  no  known 
compound  with  oxygen.  Chlorine  forms  very  unstable  com- 
pounds with  hydrogen  and  oxygen,  bromine  more  stable 
compounds,  while  iodine  forms  fairly  stable  substances  when 
combined  with  oxygen  and  hydrogen. 


BROMINE,    IODINE,    FLUORINE  143 


EXPERIMENTS  WITH  BROMINE,  IODINE,  AND  FLUORINE 

Experiment  73.  Preparation  of  Bromine  from  Potassium 
Bromide,  Manganese  Dioxide,  and  Sulphuric  Acid. — (Retort  hold- 
ing 500  cc.;  receiver;  evaporating-dish  large  enough  to  hold  the 
receiver;  potassium  bromide;  manganese  dioxide;  sulphuric 
acid  concentrated ;  ice.) 

(All  experiments  with  bromine,  iodine,  and  fluorine  must  be 
carried  out  under  the  hood.) 

Place  the  receiver  in  the  evaporating-dish  containing  ice- 
water.  Support  the  retort  on  an  iron  stand  as  in  Figure  3. 
Mix  4  grams  potassium  bromide  with  8  grams  manganese 
dioxide,  and  introduce  the  mixture  into  the  retort.  Add 
slowly  18  cc.  of  concentrated  sulphuric  acid  to  100  cc.  of  water. 
When  cold  introduce  the  mixture  into  the  retort.  Insert  the 
arm  of  the  retort  into  the  neck  of  the  receiver  and  heat  slowly 
the  contents  of  the  retort.  The  bromine  will  condense  in  the 
receiver  as  a  dark  reddish  liquid. 

Write  equations  for  above  reactions  ?     (See  page  133.) 

Experiment  74.  Solidification  of  Bromine. — (Test-tube ;  evap- 
orating-dish ;  mixture  of  finely  powdered  ice  and  salt.) 

Introduce  a  few  centimetres  of  bromine  into  a  small  test- 
tube  and  plunge  the  tube  into  a  mixture  of  finely  powdered  ice 
and  salt.  The  bromine  freezes,  forming  a  yellowish  red  solid. 
(See  page  135.) 

Experiment  75.  Preparation  of  Bromine  Water.  —  (Erlen- 
meyer  flask  holding  300  cc. ;  cork ;  bromine.) 

Introduce  into  the  flask  150  cc.  of  water ;  add  2  cc.  of  bromine. 
Cork  the  flask  and  shake  carefully  from  time  to  time.  The 
bromine  gradually  dissolves,  forming  a  yellowish  solution  known 
as  bromine  water.  Is  chlorine  soluble  in  water  ?  How  is  chlorine 
water  made  ?  (See  page  135.) 

Experiment  76.  Action  of  Bromine  on  Other  Elements.  — 
(Beaker;  test-tube;  bromine;  arsenic.) 

Into  a  small  test-tube  introduce  a  few  cubic  centimetres  of 
bromine.  Great  care  must  be  taken  not  to  allow  bromine  to 
come  in  contact  with  the  hand,  or  very  painful  wounds  will  result. 

Place  the  test-tube  in  the  beaker.  If  the  test-tube  should 
break,  the  beaker  will  catch  the  bromine  and  prevent  it  damag- 
ing the  floor  of  the  hood. 

Introduce  a  piece  of  arsenic  or  antimony  the  size  of  a  grain 
of  rice.  Vigorous  chemical  reaction  takes  place ;  the  arsenic 
or  antimony  becomes  red-hot. 


144 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


Does  chlorine  act  on  arsenic  or  antimony  ?  (See  page  134.) 
Experiment  77.  Preparation  of  Hydrobromic  Acid  by  the 
Action  of  Sulphuric  Acid  on  Potassium  Bromide.  —  (This  ex- 
periment to  be  performed  by  the  instructor.)  (Thistle-tube ;  Flor- 
ence or  balloon  flask  500  cc. ;  U-tube ;  funnel,;  beaker;  red 
phosphorus ;  potassium  bromide ;  sulphuric  acid.) 

Arrange  an  apparatus  as  shown  in  Figure  37.  Fifteen  grams 
of  potassium  bromide  are  introduced  into  the  flask  A.  The  U- 
tube  is  filled  with  a  mixture  of  red  phosphorus  and  fragments 
of  glass.  The  red  phosphorus  removes  any  bromine  from  the 


FIG.  37. 


gas,  and  the  fragments  of  glass  prevent  the  tube  from  being 
closed  by  the  red  phosphorus.  A  funnel  is  attached  to  the  end 
of  the  tube  B,  and  dips  just  beneath  the  surface  of  the  water  in 
the  beaker  C. 

Add  75  cc.  of  concentrated  sulphuric  acid  very  slowly  to 
25  cc.  water.  When  the  mixture  is  cold,  pour  into  the  flask 
through  the  thistle-tube  enough  of  the  mixture  to  cover  the 
solid  potassium  bromide.  Carefully  warm  the  flask.  Hydro- 
bromic acid  will  be  evolved,  and  will  be  readily  absorbed  by 
the  water  in  the  beaker. 

Continue  the  experiment  until  considerable  gas  has  been 
dissolved  in  the  water.  Preserve  this  solution  for  the  follow- 
ing experiment. 


BROMINE,   IODINE,   FLUORINE  145 

Write  the  equations  for  the  above  reactions  ? 

Write  the  equations  when  potassium  bromide  is  treated  with 
concentrated  sulphuric  acid  ? 

Test  the  action  of  concentrated  sulphuric  acid  on  a  crystal 
of  potassium  bromide  ?  (See  page  135.) 

Experiment  78.  Properties  of  Hydrobromic  Acid.  —  (Test- 
tubes  ;  red  litmus  paper ;  potassium  hydroxide ;  chlorine  water ; 
solution  of  hydrobromic  acid.) 

Pour  a  little  aqueous  hydrobromic  acid  into  a  test-tube ;  add 
a  drop  of  litmus  or  dip  into  the  solution  a  strip  of  blue  litmus 
paper.  Observe  what  change  takes  place  in  the  litmus.  Ex- 
plain ? 

Neutralize  another  portion  of  the  solution  of  hydrobromic 
acid  with  potassium  hydroxide.  Evaporate  the  solution  to 
dryness  on  the  water-bath.  What  is  the  substance  obtained  ? 

To  the  aqueous  solution  of  hydrobromic  acid  add  a  little 
chlorine  water.  The  solution  turns  reddish  brown.  Explain 
the  change  in  color  ?  (See  page  135.) 

Experiment  79.  Detection  of  Bromine  in  Bromides,  and  when 
in  the  Presence  of  Chlorine.  —  (Potassium  bromide ;  potassium 
chloride;  chlorine  water;  carbon  bisulphide;  silver  nitrate; 
ammonia.) 

Prepare  a  solution  of  a  few  crystals  of  potassium  bromide. 
Add  a  few  cubic  centimetres  of  carbon  bisulphide,  which  being 
heavier  than  the  solution  settles  at  once  to  the  bottom  of  the 
test-tube.  Add  chlorine  water  drop  by  drop,  shaking  vigorously 
after  the  addition  of  each  drop.  The  bromine  set  free  by  the 
chlorine  will  be  dissolved  by  the  carbon  bisulphide,  and  will 
impart  to  it  its  characteristic  red  color. 

Treat  a  dilute  solution  of  potassium  bromide  with  a  dilute 
solution  of  silver  nitrate.  White  silver  bromide  will  be  pre- 
cipitated. Equation  ?  Add  a.  few  drops  of  dilute  ammonia,  and 
the  white  precipitate  does  not  dissolve.  Compare  this  with  the 
solubility  of  silver  chloride  in  ammonia  ? 

Mix  a  gram  of  potassium  bromide  with  a  gram  of  potassium 
chloride,  and  dissolve  the  mixture.  To  a  part  of  the  mixture 
in  a  test-tube  add  a  little  carbon  bisulphide  and  then  add 
chlorine  water  drop  by  drop,  shaking  the  solution  vigorously 
after  the  addition  of  each  drop.  The  chlorine  liberates  the 
bromine,  which  gives  its  characteristic  color  to  the  carbon 
bisulphide. 

Treat  another  portion  of  the  solution  with  the  solution  of 
silver  nitrate  until  all  of  the  bromine  and  chlorine  haye  been 


146  ELEMENTS   OF   INORGANIC   CHEMISTRY 

precipitated  as  the  silver  salts.  Add  &few  drops  of  dilute  am- 
monia. The  ammonia  will  dissolve  a  part  of  the  silver  chloride, 
but  leave  undissolved  the  silver  bromide.  Filter  the  solution 
from  the  undissolved  solid,  and  add  to  the  clear  filtrate  a  little 
nitric  acid.  Silver  chloride  will  be  precipitated.  Thus  we 
detect  a  chloride  and  a  bromide  in  the  presence  of  each  other. 
(See  page  134.) 

Experiment  80.  Preparation  of  Iodine  from  Potassium 
Iodide,  Manganese  Dioxide  and  Sulphuric  Acid.  —  (Retort 
holding  300  cc.;  receiver;  potassium  iodide;  manganese  di- 
oxide; sulphuric  acid.) 

Arrange  an  apparatus  as  in  Figure  32,  omitting  the  cooling 
agent  around  the  receiver.  Mix  5  grams  of  potassium  iodide 
and  10  grains  of  finely  powdered  manganese  dioxide.  Intro- 
duce the  mixture  into  the  retort.  Add  15  cc.  of  concentrated 
sulphuric  acid  to  85  cc.  water  ;  allow  the  mixture  to  cool,  and 
introduce  into  the  retort.  Heat  gently.  Vapors  of  iodine  will 
fill  the  retort  and  condense  on  its  walls  and  in  the  receiver. 
(See  page  137.) 

Experiment  81.  Properties  of  Iodine;  Action  on  Phosphorus 
and  Zinc.  —  (Porcelain  crucible;  iodine;  phosphorus;  zinc  dust.) 

Bring  a  few  crystals  of  iodine  into  a  porcelain  crucible.    Dry 


igorous  chemical  reaction  will  quickly 
place. 

Mix  in  a  crucible  4  grams  of  iodine  which  has  been  finely 
powdered,  with  1  gram  of  zinc  dust.  Add  a  drop  or  two  of 
water.  Chemical  reaction  will  take  place.  (See  page  138.) 

Experiment  82.  Properties  of  Iodine  —  Sublimation  ;  Solu- 
bility ;  Reaction  with  Starch.  —  (500  cc.  flask  ;  carbon  bisul- 
phide; potassium  iodide;  iodine;  starch.) 

Introduce  a  few  grams  of  iodine  into  a  flask  holding  500  cc. 
Place  the  flask  on  a  sand-bath  or  an  asbestos  board,  having 
wrapped  a  piece  of  filter-paper,  moistened  with  cold  water, 
around  the  top  and  neck  of  the  flask.  The  iodine  volatilizes 
without  melting,  and  crystallizes  on  the  walls  of  the  flask  in 
beautiful  plate-like  crystals. 

While  iodine  is  only  slightly  soluble  in  water,  it  dissolves 
readily  in  carbon  bisulphide.  Shake  a  crystal  of  iodine  with  a 
few  cubic  centimetres  of  carbon  bisulphide,  and  note  the  color 
of  the  solution? 

Add  a  few  crystals  of  iodine  to  an  aqueous  solution  of  potas- 


BROMINE,    IODINE,   FLUORINE  147 

slum  iodide.  It  readily  dissolves,  forming  "tincture  of  iodine." 
Note  the  color  ? 

Prepare  a  little  starch-paste  as  follows :  Place  a  few  granules 
of  starch  in  a  mortar  and  add  enough  cold  water  to  just  moisten 
the  starch.  Grind  the  starch  to  a  paste  with  the  pestle.  Pour 
the  starch  into  a  beaker  containing  hot  (but  not  boiling)  water. 
It  is  then  ready  for  use. 

Fill  a  test-tube  about  one-third  full  with  the  starch-paste. 
To  one  drop  of  the  solution  of  iodine  in  potassium  iodide  add 
20  cc.  of  water.  Add  a  few  drops  of  this  very  dilute  solution 
of  iodine  to  the  starch-paste.  A  deep  blue  color  at  once  appears. 
This  is  a  very  sensitive  reaction  for  iodine.  (See  page  138.) 

Experiment  83.  Action  of  Sulphuric  Acid  on  Potassium 
Iodide.  —  (A  few  crystals  of  potassium  iodide;  concentrated 
sulphuric  acid.) 

Dissolve  a  few  crystals  of  potassium  iodide  in  a  little  water 
in  a  test-tube,  and  add  a  few  drops  of  concentrated  sulphuric 
acid.  We  would  expect  that  under  these  conditions  hydriodic 
acid  would  be  liberated.  The  hydriodic  acid,  however,  acts  as 
a  reducing  agent,  and  we  have  iodine  set  free  in  the  solution. 
The  hydrogen  of  the  hydriodic  acids  reacts  with  more  sul- 
phuric acid,  reducing  it  to  sulphur  dioxide,  sulphur,  or  hydro- 
gen sulphide ;  depending  upon  the  strength  of  the  solution  of 
the  acid  used,  and  upon  the  temperature.  Carefully  smell  the 
gases  escaping  from  the  test-tube  ?  (See  page  139.) 

Experiment  84.  Detection  of  Iodine.  —  (Potassium  iodide; 
carbon  bisulphide ;  chlorine  water;  silver  nitrate;  ammonia.) 

Dissolve  a  few  crystals  of  potassium  iodide  in  water ;  add  a 
solution  of  silver  nitrate.  A  yellow  precipitate  is  formed,  which 
is  insoluble  in  ammonia;  thus  differing  from  silver  chloride, 
which  is  readily  soluble  in  ammonia,  and  from  silver  bromide, 
which  is  soluble  in  a  large  volume  of  ammonia.  To  a  solution 
of  potassium  iodide  add  chlorine  water,  drop  by  drop,  having 
previously  added  a  little  carbon  bisulphide.  The  whole  is 
shaken  vigorously.  The  carbon  bisulphide  acquires  the  char- 
acteristic purplish  color  which  is  produced  by  iodine.  (See 
page  138.) 

Experiment  85.  Detection  of  Iodine  in  the  Presence  of  Chlo- 
rine and  Bromine.  —  (Potassium  iodide ;  potassium  bromide ; 
potassium  chloride  ;  carbon  bisulphide ;  chlorine  water ;  silver 
nitrate ;  ammonia.) 

Mix  1  gram  of  potassium  iodide,  1  gram  of  potassium  bro- 
mide, and  1  gram  of  potassium  chloride,  and  dissolve  the  mix- 


148  ELEMENTS   OF   INORGANIC   CHEMISTRY 

ture  in  50  cc.  of  water.  To  a  portion  of  the  solution  add  silver 
nitrate  until  all  of  the  iodine,  bromine,  and  chlorine  are  precipi- 
tated as  the  corresponding  silver  compounds.  Add  a  little 
dilute  ammonia,  filter,  and  the  chlorine  will  be  in  the  filtrate ; 
explain  ?  Add  a  little  nitric  acid  to  the  filtrate,  and  silver  chlo- 
ride will  be  reprecipitated. 

Add  an  excess  of  ammonia  to  the  mixture  of  silver  bromide 
and  iodide,  and  the  bromide  will  dissolve,  leaving  the  yellow 
silver  iodide  behind,  since  the  latter  is  insoluble  in  ammonia. 
Again  filter,  and  the  bromine  will  be  in  the  filtrate.  Add  a 
little  nitric  acid  to  the  filtrate,  and  silver  bromide  will  be  repre- 
cipitated. The  insoluble  silver  iodide  will  remain  on  the  filter- 
paper.  The  above  is  a  rough  qualitative  method  for  separating 
the  three  halogens,  — chlorine,  bromine,  and  iodine. 

A  much  simpler  method  of  detecting  iodine  in  the  presence 
of  bromine  and  chlorine  is  the  following:  To  a  little  of  the 
solution  of  the  above  mixture  add  a  few  cubic  centimetres  of 
carbon  bisulphide,  and  then  add  chlorine  water,  drop  by  drop, 
shaking  vigorously.  The  iodine  separates  first,  and  its  presence 
is  detected  by  the  purple  color  which  it  imparts  to  the  carbon 
bisulphide.  More  chlorine  water  causes  this  color  to  disappear, 
since  the  iodine  is  oxidized  to  iodic  acid,  which  is  colorless. 
Further  addition  of  chlorine  water  throws  out  the  bromine, 
which  gives  a  dull  red  color  to  the  carbon  bisulphide.  This 
color  persists  on  further  addition  of  chlorine  water,  since  bro- 
mine is  not  oxidized  by  this  reagent. 

Experiment  86.  Preparation  of  Hydrofluoric  Acid ;  Etching 
Glass.  —  (Platinum  or  lead  dish ;  powdered  fluor-spar ;  con- 

^  centrated  sulphuric 
acid ;  glass  plate ; 
paraffine.) 

Concentrated  sul- 
phuric acid  is  added 
to  finely  powdered 
fluor-spar  until  the 

^^    mass  has  the  consist- 

T*~~  ency  of  paste.    Over 

the  dish  is  placed  a 

glass  plate  covered  with  paraffine  (by  dipping  it  into  molten 
paraffine),  and  containing  a  few  lines  or  characters  drawn 
in  the  paraffine  (Fig.  38).  The  whole  is  now  set  in  a  warm 
place  and  allowed  to  remain  for  some  hours.  Where  the 
lines  were  drawn  the  paraffine  was  removed,  and  the  glass  ex- 


BROMINE,   IODINE,   FLUORINE  149 

posed  to  the  fumee  of  the  hydrofluoric  acid  set  free  by  the 
action  of  sulphuric  acid  on  calcium  fluoride.  Remove  the 
paramne  by  dissolving  it  in  warm  turpentine,  and  the  design 
will  be  permanently  etched  upon  the  glass.  Lines  are  thus 
etched  upon  measuring  apparatus.  (See  page  142.) 


PROBLEMS 

1.  Weight  of  bromine  that  can  be  obtained  from  1  kilogram 
of  potassium  bromide  ?  With  what  weight  of  arsenic  will  this 
combine  to  form  arsenic  tribromide  ? 

2  How  much  potassium  bromide  would  be  required  to  yield 
500  grams  of  hydrobromic  acid  ? 

3.  Ten  litres  of  chlorine  will  liberate  how  much  bromine 
from  potassium  bromide  ?     How  much  potassium  bromate  can 
be  formed  from  this  amount  of  bromine  ? 

4.  The  iodine  contained  in  100  grams  of  potassium  iodide 
will  form  with  phosphorus  what  weight  of  phosphorus  tri- 
iodide  ? 

5.  How  many  grams  of  potassium  iodide  can  be  obtained 
from  the  iodine  in  7  grams  of  potassium  iodate  ? 

6.  When  fluor-spar  is  decomposed  by  sulphuric  acid,  what 
amount  is  decomposed  by  12  grams  of  the  acid  ?     What  is  the 
weight  of  the  hydrofluoric  acid  liberated  ? 


CHAPTER  XV 

SULPHUR,    SELENIUM,    TELLURIUM 
SULPHUR  (At.  Wt.  =  32.06) 

Occurrence  and  Purification.  —  Sulphur  occurs  in  great  abun- 
dance in  nature  in  the  free  condition.  This  is  especially  true 
in  volcanic  regions  such  as  those  of  Italy,  Sicily,  Iceland,  etc. 

Sulphur  also  occurs  in  combination  with  a  number  of  other 
elements.  In  combination  with  oxygen  as  sulphur  dioxide, 
S02,  it  escapes  in  certain  volcanic  regions.  Combined  with 
hydrogen  as  hydrogen  sulphfde,  H2S,  it  also  issues  from  the 
earth  in  the  neighborhood  of  certain  volcanoes.  It  also  exists 
in  combination  with  a  number  of  metals  as  sulphides.  We 
have  lead  sulphide  or  galena,  PbS,  iron  sulphide  or  pyrites, 
FeS2,  antimony  sulphide  or  stibnite,  Sb2S3,  copper  iron  sulphide 
or  copper  pyrites,  Cu2Fe2S4.  The  sulphur  that  occurs  in  the  free 
condition  in  nature  does  not  all  escape  from  the  interior  of  the 
earth  in  the  form  of  sulphur,  but  is  deposited  as  the  "result  of 
the  action  of  one  sulphur  compound  on  another,  or  of  oxygen 
on  hydrogen  sulphide  :  — 


Sulphur  as  it  occurs  in  nature  contains  more  or  less  impuri- 
ties. The  sulphur  is  freed  from  these  by  fusion.  The  first 
more  or  less  crude  product  is  known  as  crude  brimstone.  The 
crude  brimstone  is  then  redistilled,  and  either  condensed  in  a 
state  of  fine  division  as  flowers  of  sulphur,  or  the  molten  mass 
poured  into  moulds  as  roll  or  stick  sulphur. 

150 


SULPHUR,    SELENIUM,    TELLURIUM  151 

Chemical  Properties  of  Sulphur.  —  Sulphur  at  ordinary  temper- 
atures is  comparatively  inert.  Indeed,  at  elevated  tempera- 
tures it  is  less  active  chemically  than  members  of  the  halogen 
group  at  ordinary  temperatures.  When  heated  in  the  presence 
of  the  air  sulphur  combines  with  oxygen,  forming  sulphur 
dioxide:-  =  SOr 


Sulphur,  however,  combines  with  many  of  the  elements, 
forming  well-defined  compounds  of  sulphur.  The  compounds 
with  the  metals,  known  as  the  sulphides,  have  characteristic 
properties  which  are  very  useful,  as  we  shall  learn,  in  qualita- 
tive analysis. 

Physical  Properties  of  Sulphur.  —  Sulphur  is  a  yellowish 
solid  at  all  ordinary  temperatures,  melting  at  118°  and  boiling 
at  448°.4.  The  solid  sulphur  is  known  in  two  forms.  If 
allowed  to  crystallize  from  a  solution  of  carbon  bisulphide,  and 
as  found  in  nature,  it  crystallizes  in  the  orthorhombic  system. 
In  this  system  the  three  axes  all  make  right  angles  with  one 
another,  and  are  all  of  unequal  lengths. 

If,  on  the  other  hand,  ordinary  flowers  of  sulphur,  roll  sul- 
phur, or  orthorhombic  sulphur  is  melted  and  allowed  to  cool 
slowly  in  a  hessian  crucible,  we  obtain  the  sulphur  in  the  form 
of  needles  which  belong  to  the  monoclinic  system.  In  the 
monoclinic  system  two  axes  are  at  right  angles  to  one  another, 
while  the  third  is  inclined.  The  three  axes  are  all  of  unequal 
lengths. 

Substances  which  can  crystallize  in  several  systems  are 
know  as  polymorphous  ;  when  they  crystallize  in  two  systems, 
as  dimorphous.  Sulphur  is,  therefore,  dimorphous.  Take  the 
two  crystalline  modifications  of  sulphur;  they  are  obviously 
analogous  to  oxygen  and  ozone. 

In  the  case  of  oxygen  and  ozone  we  have  seen  that  the 
fundamental  and  important  difference  is  in  the  amounts  of 
intrinsic  energy  present  in  the  molecules  of  the  two  sub- 
stances. We  would  naturally  ask  whether  any  such  difference 
exists  between  orthorhombic  and  monoclinic  sulphur.  This 


152  ELEMENTS   OF  INORGANIC   CHEMISTRY 

can  be  answered  very  simply  in  the  case  of  sulphur ;  indeed, 
more  simply  and  directly  than  with  oxygen  and  ozone.  It  is 
only  necessary  to  burn  equal  quantities  of  the  two  modifica- 
tions of  sulphur  in  oxygen  and  measure  the  amounts  of  heat 
liberated. 

A  considerable  difference  was  found  in  the  amounts  of  heat 
liberated  in  the  two  cases.  Thirty-two  grams,  or  a  gram-atomic 
weight,  of  orthorhombic  sulphur  when  burned  to  sulphur  di- 
oxide liberated  71,000  calories  of  heat.  An  equal  weight  of 
inonoclinic  sulphur  burned  to  sulphur  dioxide  set  free  73,300 
calories  of  heat.  The  difference,  2300  calories,  is  the  thermal 
equivalent  of  the  difference  in  the  intrinsic  energy  of  the  two 
modifications. 

Sulphur,  as  has  been  stated,  melts  at  118°.  It  first  passes 
over  into  a  thin,  light  yellow  liquid,  which,  on  further  rise  in 
temperature,  passes  through  a  remarkable  series  of  transfor- 
mations. When  heated  to  160°  the  yellow  liquid  passes  over 
into  a  reddish  brown,  viscous  mass,  which  becomes  deeper 
brown  in  color  and  more  viscous  as  the  temperature  is  raised 
to  250°.  If  the  temperature  is  still  further  raised  until  400° 
is  reached,  the  viscous  mass  becomes  a  yellow  liquid  again, 
which  boils  at  448°.4.  When  the  boiling  sulphur  is  allowed  to 
cool,  it  passes  through  these  same  changes  again,  but  in  reverse 
order. 

Hydrogen  Sulphide,  H2S.  — This  compound  occurs  in  nature 
in  the  free  condition,  as  we  have  seen.  It  escapes  from  fissures 
in  certain  localities,  and  is  dissolved  in  certain  waters,  pro- 
ducing sulphur  water. 

By  far  the  best  method  of  preparing  hydrogen  sulphide,  and 
the  one  which  is  always  employed,  is  to  treat  certain  sulphides 
with  an  acid.  The  sulphide  which  it  is  most  convenient  and 
economical  to  use  is  ferrous  sulphide,  FeS.  When  this  is  treated 
with  sulphuric  or  hydrochloric  acid,  the  following  reaction 
takes  place:  — 

FeS  +  H2S04  =  FeS04  +  H2S, 


SULPHUR,   SELENIUM,   TELLURIUM  153 

Chemical  Properties  of  Hydrogen  Sulphide.  —  The  most  im- 
portant chemical  application  of  hydrogen  sulphide  is  in  con- 
nection with  its  action  on  the  soluble  salts  of  the  heavy  metals. 
Take  as  an  example  its  action  on  solutions  of  silver  nitrate  :  — 

2  AgN03  +  H2S  =  Ag.S  +  2  HN03. 

In  such  cases  we  have  the  sulphide  of  the  metal  precipitated 
with  its  characteristic  properties. 

It  is  generally  true  that  when  hydrogen  sulphide  is  passed 
into  solutions  of  salts  of  the  heavy  metals  the  sulphide  of  the 
metal  is  precipitated.  If  this  does  not  take  place  otherwise,  it 
is  effected  by  making  the  solution  slightly  alkaline. 

The  sulphides  of  certain  metals  resemble  one  another  with 
respect  to  a  given  property.  This  enables  us  to  separate  the 
metals  into  groups  by  means  of  hydrogen  sulphide.  The 
sulphides  of  certain  metals  are  soluble  in  water. 

There  is  another  group  of  elements  whose  sulphides  dissolve 
in  dilute  acids.  These  can  be  precipitated  as  a  group,  not  by 
hydrogen  sulphide,  since  this  would  necessitate  the  formation 
of  free  acid  and  the  consequent  solution  of  the  sulphide,  but 
by  a  soluble  sulphide.  The  sulphide  employed  is  ammonium 
sulphide,  and  this  group  of  elements  is  known  as  the  ammonium 
sulphide  group. 

There  remains  a  group  of  elements  whose  sulphides  are  not 
soluble  in  cold,  dilute  acids.  Salts  of  these  metals  are  readily 
precipitated  by  hydrogen  sulphide,  since  the  acid  set  free  does 
not  dissolve  the  sulphide  when  formed. 

This  group  of  elements  is  known  as  the  hydrogen  sulphide 
group.  Here  again  individual  differences  between  the  sulphides 
are  utilized  to  separate  the  several  members. 

Hydrogen  sulphide  is  easily  oxidized  in  the  sense  of  the  fol- 
lowing equation  :  — 


Acid  Sulphides.  —  Hydrogen  sulphide,  as  we  have  seen,  has 
the  composition  H2S,  and  forms  salts  with  univalent  elements 


154  ELEMENTS   OF   INORGANIC   CHEMISTRY 

having  the  composition  M2S,  M  representing  any  univalent 
element.  If  M  represents  a  bivalent  element,  the  salt  has  the 
composition  MS. 

Hydrogen  sulphide  can  also  form  a  different  class  of  salts  in 
which  one  of  the  hydrogen  atoms  is  still  present.  With  univa- 
lent elements  these  would  have  the  general  composition  MHS, 
with  bivalent  elements  M(HS)2.  We  have  a  number  of  ex- 
amples of  these  hydrosulphides  or  acid  sulphides  as  they  are 
called.  Ammonia  forms  the  hydrosulphide  very  readily  when 
hydrogen  sulphide  is  conducted  into  aqueous  ammonia  :  — 

NH3  +  H2S  =  NH4HS. 

Indeed,  this  is  the  compound  which  is  formed  when  aqueous 
ammonia  is  saturated  with  hydrogen  sulphide. 

Dissociation  of  Hydrogen  Sulphide.  —  The  hydrosulphides  or 
acid  sulphides  are,  as  we  have  seen,  salts  of  a  monobasic  acid. 
The  sulphides,  on  the  other  hand,  are  salts  of  a  dibasic  acid, 
i.e.  one  which  combines  with  two  univalent  ions.  How  can  we 
explain  these  facts  on  the  basis  of  the  dissociation  theory  ? 

There  are  two  different  ways  in  which  hydrogen  sulphide 
can  dissociate.  These  are  the  following  :  — 

H2S  =  H,HS,  (1) 

H,H,S.  (2) 


When  the  compound  dissociates  as  in  equation  (2),  it  is 
obviously  dibasic;  when  the  dissociation  takes  place  as  in 
equation  (1),  it  is  monobasic. 

Physical  Properties  of  Hydrogen  Sulphide.  —  Hydrogen  sul- 
phide is  a  colorless  gas,  with  an  extremely  disagreeable  odor, 
suggesting  decomposing  organic  matter.  When  taken  into 
the  lungs  in  any  great  quantity,  it  is  quite  poisonous.  For 
this  reason  and  on  account  of  its  disgusting  odor  it  should 
never  be  allowed  to  escape  into  the  air  of  a  laboratory.  Since 
it  is  so  frequently  used  in  connection  with  qualitative  analysis, 
a  separate  room  is  attached  to  every  well-equipped  chemical 


SULPHUR,    SELENIUM,   TELLURIUM  155 

laboratory  in  which  the  gas  is  generated  and  used.     This  is 
known  as  the  sulphuretted  hydrogen  room. 

Hydrogen  sulphide  dissolves  in  water  to  the  extent  of  from 
2^  to  3  volumes  of  the  gas  in  one  volume  of  water. 

\ 

COMPOUNDS  OF   SULPHUR  WITH  OXYGEN    AND 
HYDROGEN 

Sulphur  Dioxide,  S02.  —  The  simplest  compound  of  sulphur 
and  oxygen  is  sulphur  dioxide,  having  the  composition  S02.  It 
is  formed  by  the  direct  combination  of  the  two  elements  when 
sulphur  is  burned  in  oxygen :  — 

S  +  02  =  S02. 

A  more  convenient  method  of  preparing  sulphur  dioxide  is 
by  the  action  of  strong  acids  on  sulphites,  or  acjd  sulphites. 
Normal  sodium  sulphite  has  the  composition  Na2SD3.  When 
this  is  treated  with  sulphuric  acid  the  following  reaction  takes 
place :  — 

Na2S03  +  H2S04  =  Na2S04  +  H20  +  S02. 

Another  very  convenient  method  for  preparing  sulphur 
dioxide  is  by  the  action  of  sulphuric  acid  on  metallic  copper :  — 

Cu  +  2  H2S04  =  CuS04  +  2  H20  +  S02. 

It  is  a  colorless  gas  with  very  penetrating  odor,  and  a  taste 
which  persists  for  an  unusual  time.  Its  odor  and  taste  are 
characteristic  of  a  burning  sulphur  match,  with  which  every 
one  is  familiar.  It  readily  takes  up  oxygen  and  is,  therefore, 
an  excellent  reducing  agent. 

Sulphurous  Acid,  H2S03.  —  The  acid  formed  when  sulphur 
dioxide  is  dissolved  in  water  yields  salts  having  the  composi- 
tion M2S03,  where  M  is  an  univalent  metal.  The  acid  must 
therefore  have  the  composition  H2S03,  and  be  a  dibasic  acid. 
It  can  also  form  acid  sulphites  of  the  composition  MHS03. 


156  ELEMENTS  OF  INORGANIC   CHEMISTRY 

It  has  been  impossible  to  isolate  the  acid  H2S03,  since  it 
breaks  down  so  readily  into  water  and  sulphur  dioxide :  — 

H2S03  =  H20  +  S02. 

Dissociation  of  Sulphurous  Acid.  —  Sulphurous  acid  is  a 
weak  acid.  Being  a  dibasic  acid,  it  can  dissociate  in  two  ways. 

H2S03  =  H,  HS03 

is  the  first  stage  in  the  dissociation.  The  second  stage  in  the 
dissociation,  produced  only  in  very  dilute  solutions  since  sul- 
phurous acid  is  a  weak  acid,  would  then  be  represented  thus :  — 

HS03  =  H,  S03. 

Sulphur  Trioxide,  S03.  —  Sulphur  trioxide,  S03,  is  formed  by 
gently  heating  fuming  sulphuric  acid.  The  latter  is  really 
a  solution  of  sulphur  trioxide  in  sulphuric  acid,  usually  having 
approximately  the  composition  H2S207.  When  this  is  heated, 
it  decomposes  thus  :  — 

H2S207  =  H2S04+S03. 

When  sulphur  dioxide  and  oxygen  are  heated  together,  they 
combine  and  form  sulphur  trioxide,  but  only  in  very  small 
quantity.  If,  however,  the  mixture  of  the  two  gases  is  heated 
in  the  presence  of  certain  substances,  they  combine  readily, 
forming  sulphur  trioxide.  Such  a  substance  is  finely  divided 
platinum.  If  the  two  gases  are  passed  through  a  tube  con- 
taining platinum  sponge,  and  the  tube  heated,  they  combine 
very  readily. 

This  is  distinctively  a  catalytic  reaction,  i.e.  a  reaction 
effected  by  a  substance  which  does  not  enter  into  the  reaction, 
but  acts  simply  by  contact.  The  platinum  in  this  case  does 
not  enter  at  all  into  the  reaction,  and  a  very  small  amount 
of  platinum  is  capable  of  effecting  a  large  amount  of  the 
reaction.  This  method  of  preparing  sulphur  trioxide  has  been 
found  to  be  so  efficient  and  economical  that  it  is  rapidly  sup- 
planting all  others,  and  in  the  near  future  will  probably  be 
used  almost  exclusively  for  preparing  this  substance. 


SULPHUR,    SELENIUM,   TELLURIUM  157 

Sulphur  trioxide  is  a  powerful  oxidizing  agent,  readily  giving 
up  one  of  its  oxygen  atoms  and  passing  over  into  sulphur 
dioxide.  It  has  an  unusual  attraction  for  water,  combining 
with  it  at  once  on  mere  contact,  and  even  causing  the  hydrogen 
and  oxygen  in  organic  compounds  to  combine  and  form  water, 
with  which  it  instantly  combines. 

Sulphur  trioxide  is  a  transparent,  mobile  liquid,  which  boils 
at  46°. 2.  It  can  be  cooled  to  zero  without  solidifying.  When 
further  cooled  it  forms  a  white  solid,  which  melts  at  14°.8. 

Sulphuric  Acid,  H2S04.  —  Sulphuric  acid  occurs  in  the  free 
condition  in  small  quantity  in  certain  waters  on  the  surface 
of  the  earth.  It  is  prepared  now  very  largely  by  the  method 
described  above, — by  the  action  of  water  on  sulphur  trioxide,  — 

S03  +  H20  =  H2S04, 

the  trioxide  being  easily  formed  by  the  direct  union  of  sulphur 
dioxide  and  oxygen  in  the  presence  of  finely  divided  platinum. 
At  an  earlier  date  a  method  was  employed  for  effecting  the 
oxidation  of  sulphur  dioxide,  which  even  to-day  finds  consider- 
able application.  This  method  is  based  upon  the  oxidation  of 
sulphur  dioxide  in  the  presence  of  water,  nitric  acid,  and  air. 

The  sulphur  dioxide  produced  by  roasting  sulphides,  or  heat- 
ing sulphur  in  contact  with  an  abundant  supply  of  air,  is  con- 
ducted through  a  tower  known  as  the  Glover  tower.  This  is 
filled  with  fire-brick,  over  which  dilute  sulphuric  acid  trickles. 
The  gas  is  cooled,  and,  at  the  same  time,  the  dilute  sulphuric 
acid  loses  water  and  is  concentrated. 

Concentrated  sulphuric  acid,  containing  the  oxides  of  nitro- 
gen, is  brought  in  contact  with  the  more  dilute  sulphuric  acid, 
when  it  sets  free  these  oxides  which  now  mix  with  the  sulphur 
dioxide.  Nitric  acid  is  then  introduced  in  the  form  of  vapor, 
and  also  water-vapor.  The  mixture  of  gases  containing  sul- 
phur dioxide,  nitric  acid,  oxides  of  nitrogen,  and  water-vapor 
are  conducted  into  a  chamber  lined  with  lead,  —  the  so-called 
leaden  chamber.  In  these  chambers,  of  which  there  are  a  series, 
the  oxidation  of  the  sulphur  dioxide  to  sulphur  trioxide  takes 


158  ELEMENTS   OF  INORGANIC   CHEMISTRY 

place,  and  the  union  of  the  sulphur  trioxide  with  water,  forming 
sulphuric  acid. 

In  order  to  avoid  loss  of  nitric  acid  and  oxides  of  nitrogen, 
which  are  expensive,  the  sulphuric  acid  formed  in  the  leaden 
chambers  is  passed  through  another  tower,  known  as  the  Gay- 
Lassac  tower.  This  tower  contains  pieces  of  coke  over  which 
concentrated  sulphuric  acid  flows.  The  concentrated  acid  takes 
up  the  oxides  of  nitrogen.  This  acid  is  then  passed  into  the 
Glover  tower,  where  it  mixes  with  the  more  dilute  sulphuric 
acid  and  gives  up  the  oxides  of  nitrogen,  which  then  oxidize 
more  sulphur  dioxide  to  the  trioxide,  and  the  process  is  thus  a 
continuous  one,  —  a  small  amount  of  nitric  acid  serving  to 
oxidize  a  large  amount  of  sulphur  dioxide. 

The  chemical  reactions  which  take  place  in  the  manufacture 
of  sulphuric  acid  by  the  above  method,  as  far  as  they  are 
known,  are  the  following :  — 

Nitric  acid  acts  upon  sulphur  dioxide  in  the  presence  of 
water  as  follows  :  — 

2  S02  +  2  HN03  +  H20  =  2  H2S04  +  N203. 

When  the  sesquioxide  of  nitrogen,  N203,  reacts  with  water- 
vapor  it  forms  nitrous  acid :  — 

N203  +  H20  =  2  HN02. 

When  nitrous  acid  reacts  with  sulphur  dioxide  in  the  pres- 
ence of  the  oxygen  of  the  air,  we  have  :  — 

2  S02  +  2  HN02  +  02  =  2  N02.S03H. 

The  compound  N02.S03H  is  known  as  nitrosyl-sulplmric  add, 
or  nitrosulphonic  acid.  When  this  is  treated  with  water,  the 
following  decomposition  takes  place  :  — 

N02.  S03H  +  H20  =  H2S04  +  HN02. 

The  HN02,  or  its  anhydride,  N203,  then  reacts  with  more  sul- 
phur dioxide  and  oxygen  and  forms  again  N02,  S03H,  which 
then  decomposes  with  water-vapor  in  the  sense  of  the  last  equa- 
tion, and  the  process  is  continuous;  the  nitrous  acid  or  sesqui- 


SULPHUR,   SELENIUM,   TELLURIUM  159 

oxides  of  nitrogen  being  collected  in  the  Gay-Lussac  tower,  as 
we  have  seen. 

The  acid  obtained  from  the  leaden  chambers  is  known  as 
chamber  acid.  It  contains  about  65  per  cent  of  the  com- 
pound H2S04.  In  order  to  further  concentrate  this  acid  it  is 
allowed  to  flow  through  hot,  shallow,  lead  pans,  and  when  the 
acid  has  become  sufficiently  concentrated  to  act  chemically 
upon  the  lead,  it  is  transferred  to  a  platinum  vessel  and  more 
of  the  water  distilled  off.  The  acid  can  be  still  further  con- 
centrated in  vessels  of  platinum. 

Chemical  Properties  of  Sulphuric  Acid.  —  One  of  the  most 
characteristic  properties  of  sulphuric  acid  is  its  power  to  take 
up  water  and  combine  with  it.  For  this  reason  it  is  an  excel- 
lent drying  agent,  readily  taking  water  from  other  substances. 
When  we  wish  to  dry  a  gas  which  contains  water-vapor,  the 
best  method  with  one  exception  is  to  allow  the  gas  to  stream 
slowly  in  fine  bubbles  through  concentrated  sulphuric  acid. 
Its  power  to  combine  with  water  is  the  key  to  many  of  the 
reactions  which  concentrated  sulphuric  acid  can  effect.  When 
brought  in  contact  with  many  organic  substances,  it  causes  the 
hydrogen  and  oxygen  to  combine  and  form  water  with  which 
it  itself  combines.  This  is  the  explanation  of  the  charring  of 
wood  and  similar  substances  effected  by  the  concentrated 
acid. 

Sulphuric  acid  has  the  power  of  driving  more  volatile  acids 
out  of  their  salts,  combining  with  the  metal  of  the  salt.  Thus, 
when  a  dry  chloride  or  nitrate  is  treated  with  sulphuric  acid, 
the  hydrochloric  or  nitric  acid  is  driven  out  and  the  sulphate 
of  the  metal  is  formed :  — 

2  NaCl  +  H2S04  =  Na2S04  +  2  HC1 ; 
2  NaN03  -f  H2S04  =  Na2S04  +  2  HN03. 

This  might  seem  to  argue  that  sulphuric  acid  was  a  stronger 
acid  than  either  hydrochloric  or  nitric.  Such,  however,  is  not 
the  case.  Hydrochloric  and  nitric  acids  being  volatile  are  dis- 
placed by  much  weaker,  non-volatile  acids,  in  accordance  with 


160  ELEMENTS   OF   INORGANIC   CHEMISTRY 

the  general  principle  that  whenever  a  volatile  compound  can  be 
formed  and  escape  from  the  field  of  action,  it  is  formed. 

Sulphuric  acid  combines  with  most  of  the  metals  or  base- 
forming  elements,  forming  sulphates.  In  all  of  these  com- 
pounds the  sulphuric  acid  is  dibasic,  combining  with  two 
atoms  of  a  univalent  metal,  with  one  atom  of  a  bivalent  metal, 
and  so  on.  The  sulphates  are  stable,  well-crystallized  com- 
pounds. The  sulphates  of  the  heavy  metals  are  generally  only 
slightly  soluble  in  water.  The  insolubility  of  its  barium  salt 
even  in  dilute  acids  furnishes  us  with  a  means  of  detecting 
sulphuric  acid  and  determining  it  quantitatively. 

Dissociation  of  Sulphuric  Acid.  —  Sulphuric  acid  is  a  typical 
dibasic  acid,  forming  two  well-defined  classes  of  salts  —  the 
normal  sulphates  and  the  acid  sulphates.  The  former  have  the 
composition  M2S04  and  the  latter  MHS04. 

Like  dibasic  acids  in  general  sulphuric  acid  dissociates  in 
two  stages.  At  first  it  breaks  down,  thus  :  — 

H2S04  =  H,  HS04. 

When  the  dilution  of  the  solution  is  increased,  i.e.  when 
more  water  is  added,  the  ion  HS04  dissociates  thus :  — 

HS04  =  H,  S04. 

Sulphuric  acid  is  a  strong  dibasic  acid,  and,  therefore,  the 
ion  HSO4  dissociates  into  H  and  S04  before  any  very  great 
dilution  is  reached. 

Physical  Properties  of  Sulphuric  Acid,  —  Sulphuric  acid,  or 
oil  of  vitriol,  is,  as  its  name  implies,  a  liquid.  It  is  a  thick, 
oily  liquid,  with  a  specific  gravity  of  1.84.  It  boils  at  338°. 

Scientific  and  Technical  Uses  of  Sulphuric  Acid.  —  Sulphuric 
acid  is  used  very  frequently  in  the  scientific  laboratory,  and  far 
more  frequently  in  technical  processes  than  any  other  acid. 
In  scientific  operations  it  is  used  as  a  dehydrating  agent,  as  a 
drying  agent  to  liberate  volatile  acids  from  their  compounds, 
and  in  many  other  processes.  In  the  arts  it  is  used  on  every 


SULPHUR,    SELENIUM,    TELLURIUM  161 

hand,  and  crude  sulphuric  acid  is  manufactured  by  the  hundreds 
of  thousands  of  tons  annually.  It  is  used  to  render  normal 
phosphates  soluble  in  water  by  converting  them  into  acid  phos- 
phates, which  can  be  assimilated  by  plants.  These  are  the 
basis  of  most  of  the  commercial  fertilizers.  It  is  also  used  in 
connection  with  the  manufacture  of  chlorine  from  sodium 
chloride,  and  in  the  preparation  of  soda. 

Sulphuric  acid  is  at  present  extensively  used  in  connection 
with  the  generation  of  electrical  energy  in  accumulators  or 
storage  cells.  In  such  cells  the  electrodes  are  plates  of  lead, 
and  the  electrolyte  dilute  sulphuric  acid. 

Other  Compounds  of  Sulphur  with  Oxygen  and  Hydrogen.  — 
The  two  acids  already  considered,  sulphurous  and  sulphuric, 
are  the  most  important  compounds  of  sulphur  with  oxygen  and 
hydrogen.  Several  other  compounds  of  sulphur  with  oxygen 
and  hydrogen  are  known. 

The  sodium  salt  of  thiosulphuric  acid  is  of  interest  in  connec- 
tion with  photography.  It  is  frequently  known  in  the  arts 
as  hyposulphite  or  simply  as  hypo.  Its  solution  dissolves  the 
halogen  salts  of  silver,  and  it  is,  therefore,  used  for  "  fixing " 
photographs.  It  is  formed  by  boiling  sodium  sulphite  with 
sulphur :  — 


SELENIUM  (At.  Wt.  =  79.2) 

Selenium.  —  Selenium  occurs  in  the  same  general  associations 
as  sulphur,  and  frequently  along  with  sulphur.  It  occurs  in 
combination  with  silver  and  copper  as  definite  minerals.  It  is 
frequently  found  in  the  dust  of  flues  where  sulphides  are 
roasted,  or  in  the  chambers  in  the  manufacture  of  sulphuric 
acid.  Selenium  combines  directly  with  hydrogen,  forming  the 
compound  H2Se,  —  hydrogen  selenide,  —  which  is  strictly  anal- 
ogous to  hydrogen  sulphide. 

Selenium  combines  with  oxygen,  forming  selenium  dioxide, 
the  analogue  of  sulphur  dioxide. 


162  ELEMENTS   OF  INORGANIC  CHEMISTRY 

Selenious  acid,  formed  by  the  union  of  selenium  dioxide  with 

\V  r\  1~  P  7*  '-" 

Se02+H20=H2Se03, 

resembles  in  many  respects  sulphurous  acid. 

While  the  compound,  selenium  trioxide,  is  not  known,  the 
acid  corresponding  to  this  anhydride  is  known.  When 
selenium  is  treated  with  strong  oxidizing  agents,  selenic  acid  or 
its  salt  is  formed. 

TELLURIUM  (At.  Wt.  =  127.6) 

Tellurium.  —  Tellurium  is  a  much  rarer  element  than 
selenium.  Tellurium  combines  with  hydrogen,  forming  hy- 
drogen telluride,  having  the  composition  H2Te.  This  is  anal- 
ogous to  hydrogen  sulphide  and  hydrogen  selenide.  It  is 
a  gas  with  a  very  disagreeable  odor,  like  the  former  com- 
pounds. 

Tellurium  combines  with  oxygen,  forming  the  compounds 
TeO,  Te02,  and  Te03.  The  last  two  are  analogous  to  sulphur 
dioxide  and  sulphur  trioxide,  while  the  first  has  no  analogue 
among  the  sulphur  compounds. 

Tellurium,  however,  forms  two  acids  with  hydrogen  and 
oxygen.  These  are  tellurians  and  telluric  acids,  having  the 
compositions  respectively  H2Te03  and  H2Te04. 


EXPERIMENTS  WITH  SULPHUR,  SELENIUM,  AND 
TELLURIUM 

Experiment  87.  Fusion  and  Distillation  of  Sulphur.  —  (Re- 
tort ;  roll  sulphur ;  beaker  filled  with  water.) 

Arrange  the  retort  as  shown  in  Figure  39.  Add  sulphur 
to  a  retort  holding  150  cc,  until  it  is  about  one-fourth  full. 
Heat  slowly.  The  sulphur  melts,  forming  a  light  yellow  liquid. 
Continue  to  raise  the  temperature,  and  the  liquid  will  become 
darker  in  color  and  more  viscous.  Heat  to  a  still  higher  tem- 
perature, and  the  sulphur  will  boil  and  pass  out  of  the  retort 
in  the  form  of  vapor,  or  condense  as  a  liquid  in  the  neck  of 
the  retort. 


SULPHUR,  SELENIUM,   TELLURIUM 


163 


If  the  liquid  sulphur  is  allowed  to  run  into  a  beaker  contain- 
ing cold  water,  it  forms  a  tough,  plastic  solid,  which  is  more 
or  less  elastic.     This  differs 
fundamentally  from  ordinary 
solid  sulphur.  (See  page  151.) 

Experiment  88.  Different 
Crystal  Forms  of  Sulphur.  — 
(Erlenmeyer flask  250  cc. ;  roll 
sulphur ;  carbon  bisulphide  ; 
porcelain  or  sand  crucible.) 

Add  400  cc.  of  carbon  bisul- 
phide to  the  Erlenmeyer 
flask,  and  then  add  25  grams 
of  sulphur.  Cork  the  flask 
and  shake  vigorously  from 
time  to  time  for  an  hour  or 
two.  Allow  the  flask  to  stand 
quietly  until  the  solution  be- 
comes clear.  Pour  off  the 
clear  solution  into  an  evap- 
orating-dish  and  allow  the 
carbon  bisulphide  to  evapo- 
rate spontaneously.  Crystals 
of  sulphur  will  be  obtained 
having  a  given  crystal  form  ? 
These  crystals  are  stable,  i.e. 
do  not  undergo  any  change 
if  allowed  to  stand  for  any  Fio.  39. 

length  of  time. 

A  porcelain  or  sand  crucible  holding  not  less  than  100  cc. 
is  filled  half  full  of  roll  sulphur.  The  sulphur  is  completely 
melted  and  the  crucible  then  covered  and  placed  in  sand  in 
which  it  cools  slowly.  When  a  crust  has  formed  on  the  sur- 
face, this  is  punctured  and  the  molten  sulphur  poured  out  of 
the  crucible.  The  walls  of  the  crucible  will  be  covered  with 
fine  needles,  which  may  extend  well  into  the  crucible.  Do  not 
touch  the  needles  with  the  hand,  but  remove  a  few  of  them 
carefully  with  a  porcelain  spatula.  Compare  these  with  sul- 
phur crystals  obtained  from  the  carbon  bisulphide.  What 
difference  in  color,  transparency,  and  form  do  you  note  ? 

Set  some  of  the  crystals  aside  for  a  day  or  two  and  examine 
them  again.  They  are  markedly  changed.  What  changes  do 
you  note  ? 


164 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


This  form  of  sulphur,  known  as  monoclinic  sulphur,  is  un- 
stable at  ordinary  temperatures  and  passes  over  into  the 
orthorhombic  form,  which  is  the  form  obtained  from  the 
carbon  bisulphide.  (See  page  151.) 

Experiment  89.  Reaction  of  Sulphur  with  Metals.  —  (Zinc 
dust;  copper  foil  or  copper  gauze;  iron  powder;  silver  coin; 
sulphur.) 

Mix  1  gram  of  sulphur  with  2  grams  of  zinc  powder,  and 
introduce  the  mixture  into  a  hot  porcelain  crucible  or  metal 
dish.  What  evidence  have  you  of  chemical  action  ?  Treat 
the  product  with  hydrochloric  acid  and  smell  the  escaping  gas. 
Treat  some  of  the  uiiheated  mixture  with  hydrochloric  acid  and 
smell  the  gas  which  escapes.  Do  the  two  have  the  same  odor  ? 
Explain  ? 

Melt  a  few  grams  of  sulphur  in  a  large  test-tube  and  heat 
the  sulphur  to  boiling.  Introduce  into  the  boiling  sulphur 
small  pieces  of  copper  foil  or  copper  gauze.  What 
evidence  is  there  of  chemical  action  ?  Examine 
the  copper  strip  when  cold.  Has  any  change  oc- 
curred ? 

Mix  20  grams  of  sulphur  with  10  grams  of  iron 
powder.  Place  some  of  the  mixture  in  a  large 
hard-glass  tube  and  test  carefully.  Is  there  any 
evidence  of  chemical  action  ? 

Treat  the  product  with  an  acid  and  smell  the 

escaping  gases. 

Put  sulphur  on  a 
moist  silver  coin. 
Does  any  change  take 
place  ?  (See  page 
151.) 

Experiment  90. 
Preparation  of  Hydro- 
gen Sulphide. — (Flor- 
ence or  balloon  flask 
holding  500  cc. ;  this- 
tle -  tube  ;  two  -  hole 
stopper;  glass  tubing; 

FIG  ^  glass  cylinder  closed 

by   a   two-hole   stop- 
per ;  ferrous  sulphide ;  dilute  hydrochloric  acid.) 

Set  up  the  apparatus  as  in  Figure  40.  Cover  the  bottom  of  the 
flask  A  with  pieces  of  ferrous  sulphide  and  add  dilute  hydro- 


SULPHUR,    SELENIUM,   TELLURIUM 


165 


chloric  acid  through  the  thistle-tube  B.  The  gas  is  passed 
through  water  in  the  cylinder  C,  and  can  then  be  passed  into 
water  giving  a  solution  of  hydrogen  sulphide,  or  collected  in 
cylinders  by  displacement  of  air.  Since  hydrogen  sulphide  is 
much  heavier  than  air,  the  cylinder  which  is  to  receive  the  gas 
must  be  placed  with  its  mouth  upwards. 

On  account  of  the  disagreeable  odor  and  poisonous  proper- 
ties of  hydrogen  sulphide,  all  experiments  with  this  gas  must 
be  performed  under  the  hood.  (See  page  152.) 

Experiment  91.  Hydrogen  Sulphide  burns  in  Contact  with 
the  Air.  —  (Tall  cylinder  or  bottle  filled  with  hydrogen  sul- 
phide ;  a  taper.) 

Lower  a  lighted  taper  slowly  into  a  cylinder  or  bottle  filled 
with  hydrogen  sulphide.  The  taper  will  be  extin- 
.  guished,  showing  that  hydrogen  sulphide  will  not 
support  combustion.  The  hydrogen  sulphide,  how- 
ever, will  take  fire  and  burn  at  the  mouth  of  the 
vessel.  The  combustion  will  not  be  complete,  and 
sulphur  will  be  deposited  on  the  walls  of  the  vessel. 
(See  page  153.) 


FIG.  41. 

Experiment  92.    Action  of  Hydrogen  Sulphide  on  Salts  of 

Metals.  —  (Balloon  flask  with  a  long  entrance-tube ;  four  cylin- 
ders each  fitted  with  a  two-hole  rubber  stopper;  Erlenmeyer 
flask  holding  250  cc. ;  glass  tubing ;  ferrous  sulphide  ;  dilute 
hydrochloric  acid;  zinc  sulphate;  lead  nitrate;  cadmium 
chloride ;  antimonious  chloride.) 

Set  up  the  apparatus  as  in  Figure  41.  If  a  Kipp  apparatus 
for  generating  hydrogen  is  available,  it  may  be  used  instead 
of  the  balloon  flask.  The  tube  B  must  be  about  40  cm.  long. 


166 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


Introduce  pieces  of  ferrous  sulphide  into  A ;  a  dilute  solution 
of  zinc  sulphate  into  (7;  of  lead  nitrate  into  D;  of  cadmium 
chloride  into  E ;  and  of  antimonious  chloride  to  which  a  little 
tartaric  acid  has  been  added  into  F.  An  alkali  is  added  to  G 
to  catch  any  gas  which  may  pass  through  all  of  the  cylinders, 
and  prevent  it  from  escaping  into  the  air  of  the  hood. 

Pour  hydrochloric  acid  through 
B  into  A.  The  gas,  hydrogen  sul- 
phide, will  be  evolved  and  pass 
through  the  solutions.  As  soon 
as  all  of  the  air  is  driven  out,  a 
white  precipitate  of  zinc  sulphide 
will  begin  to  form  in  C ;  a  black 
precipitate  of  lead  sulphide  will 
form  in  D.  The  yellow  precipi- 
tate in  E  is  cadmium  sulphide; 
while  orange-red  antimony  sul- 
phide will  be  precipitated  in  F. 

These  are  some  of  the  most  char- 
acteristic reactions  of  hydrogen 
sulphide.  Nearly  all  of  the  heavy 
metals,  as  we  shall  see,  are  pre- 
cipitated from  solutions  of  their 
salts  by  hydrogen  sulphide  di- 
rectly, or  by  the  addition  of  am- 
mcnia  after  the  solution  has  been 
saturated  with  hydrogen  sul- 
phide. (See  page  135.) 

Experiment  93.  Preparation 
of  Sulphur  Dioxide  by  burning 
Sulphur  in  Oxygen.  —  (Large 
glass  tubing;  porcelain  boat;  gasometer  filled  with  oxygen; 
sulphur.)  The  gasometer  (Fig.  42)  is  filled  with  oxygen  as 
follows  :  The  stop-cock  C  is  opened.  The  reservoir  D  is  filled 
with  water  and  the  stop-cock  B  opened.  After  filling  the  gas- 
ometer with  water  the  stop-cocks  C  and  B  are  closed  and  the 
cap  A  removed.  The  discharge-tube  from  an  oxygen  generat- 
ing apparatus  is  introduced  into  A,  and  as  the  oxygen  enters 
the  gasometer,  water  escapes  through  A. 

A  hard-glass  tube  20  cm.  long  is  closed  at  both  ends  by 
stoppers,  through  which  pass  small  glass  tubes  as  in  Fig- 
ure 43.  A  plug  of  asbestos  is  introduced  into  one  end  of 
the  tube,  and  oxygen  passed  through  the  tube  until  all  the 


FIG.  42. 


SULPHUR,   SELENIUM,   TELLURIUM 


16T 


air  has  been  driven  out.  The  porcelain 
boat,  containing  a  gram  or  two  of  sulphur, 
is  now  introduced  into  the  other  end  of 
the  glass  tube  and  ignited.  The  cork  is 
replaced  and  the  stream  of  oxygen  contin- 
ued. The  sulphur  burns  in  the  oxygen, 
forming  sulphur  dioxide.  The  gas  passes 
out  of  the  tube,  any  sulphur  vapors  being 
held  back  by  the  asbestos. 

The  gas  can  be  collected  in  cylinders  by 
displacement  of  air,  not  over  water,  since 
it  is  so  very  soluble  in  water,  or  it  may  be 
passed  into  water,  and  gives  a  solution  of 
sulphurous  acid. 

Test  the  reaction  of  the  aqueous  solution 
towards  litmus  ?  (See  page  155.) 


FIG.  43. 


FIG,  44. 


Experiment  94.  Preparation  of  Sul- 
phur Dioxide  by  the  Action  of  Sulphuric 
Acid  on  Acid  Sodium  Sulphite.  —  (Flask 
holding  500  cc. ;  acid  sodium  sulphite ; 
a  mixture  of  one-half  sulphuric  acid  and 
one-half  water.) 

Pour  25  grams  of  acid  sodium  sul- 
phite into  the  flask  (Fig.  44)  and  slowly 
drop  the  sulphuric  acid  on  to  the  sulphite 
from  the  dropping-funnel.  A  steady 
stream  of  sulphur  dioxide  will  be  evolved. 
(See  page  155.) 

Experiment  95.  Preparation  of  Sul- 
phur Dioxide  by  the  Action  of  Sulphuric 
Acid  on  Copper.  —  (Flask  holding  500  cc. ; 
safety  -  tube ;  discharge  -  tube ;  funnel ; 
glass  vessel.) 

Arrange  the  apparatus  as  in  Figure 
45.  Introduce  some  copper  turnings 
into  the  flask  A,  and  cover  these  with 
concentrated  sulphuric  acid.  Warm 
gently,  when  sulphur  dioxide  will  escape 
and  can  be  dissolved  in  water.  Sulphur 


168 


ELEMENTS  OF  INORGANIC   CHEMISTRY 


dioxide  is  very  soluble  in  water,  and  therefore  the  funnel  is 
used  to  prevent  the  return  of  the  water  into  the  flask  A.  (See 
page  155.) 


FIG.  45. 

Experiment  96.  Liquefaction  of  Sulphur  Dioxide.  —  (Flask 
holding  500  cc. ;  dropping-f unnel ;  gas  wash-bottle;  U-tube; 
evaporating-dish  ;  acid  sodium  sulphite ;  sulphuric  acid ;  salt 
and  ice.)  (  To  be  performed  by  the  instructor.) 

Arrange  the  apparatus  as  in  Figure  46.  The  acid  sulphite  is 
introduced  into  the  flask  A,  and  the  mixture  of  equal  parts  of 
water  and  sulphuric  acid  dropped  slowly  on  the  salt.  The  gas 
is  passed  through  sulphuric  acid  in  the  cylinder  B,  and  then 
through  the  U-tube  C,  surrounded  by  ice  and  salt.  The  gas 
condenses  to  a  liquid  in  the  U-tube. 

When  liquid  sulphur  dioxide  evaporates,  which  it  does  readily 
at  ordinary  temperatures,  a  low  temperature  is  produced.  Mix 
in  a  test-tube  a  few  cubic  centimetres  of  liquid  sulphur  dioxide 
with  an  equal  volume  of  water.  The  water  will  be  frozen. 

Experiment  97.  Bleaching  Action  of  Moist  Sulphur  Dioxide. 
—  (Bell-jar;  tripod j  crucible;  flowers;  some  colored  fabric j 
sulphur.) 


SULPHUR,   SELENIUM,   TELLURIUM 


169 


Place  a  few  grams  of  sulphur  in  the  porcelain  crucible.     Place 
some  flowers  and  a  moist  piece  of  calico  fabric  on  the  tripod 


FIG.  46. 

(Fig.  47).     Ignite  the  sulphur  and  cover  the  whole  with  a  bell- 
jar  resting  on  a  glass  plate.     In  a  short 
time  the  flowers   and  sulphur  will  be 
decolorized.     (See  page  155.) 

Experiment  98.  Reducing  Action  of 
Sulphurous  Acid.  —  (Potassium  perman- 
ganate; potassium  dichromate;  alcoholic 
solution  of  iodine ;  aqueous  solution  of 
sulphur  dioxide.) 

To  an  aqueous  solution  of  potassium 
permanganate  add,  drop  by  drop,  the 
sulphurous  acid.  The  beautiful  color  of 
the  permanganate  will  finally  entirely 

lisappear.     Reverse  the  procedure,  and 

Irop   the  permanganate    into  the    sul- 

)hurous    acid.      The   disappearance   of 
color  can  be  readily  seen. 


In  a  similar  manner  add  sulphurous 
acid  to  a  solution  of  potassium  dichro- 


FIG.  47. 


170  ELEMENTS  OF  INORGANIC  CHEMISTRY 

mate.  The  red  color  will  disappear  and  a  green  color  will  be 
produced.  Add  sulphurous  acid  to  an  alcoholic  solution  of 
iodine.  It  will  become  colorless. 

In  all  of  the  above  reactions  sulphurous  acid  acts  as  a  reduc- 
ing agent,  it  being  oxidized  to  sulphuric  acid.  (See  page  155.) 

Experiment  99.  Properties  of  Sulphuric  Acid.  —  (Pine  splin- 
ter ;  cane  sugar ;  litmus ;  methyl  orange  j  dilute  sulphuric  acid ; 
concentrated  sulphuric  acid.) 

The  power  of  sulphuric  acid  to  combine .  with  water  effects  a 
number  of  chemical  reactions.  Introduce  a  pine  splinter  into 
concentrated  sulphuric  acid.  It  is  quickly  charred.  The  sul- 
phuric acid  has  caused  the  hydrogen  and  oxygen  in  the  wood 
to  combine  and  form  water,  which  it  takes  up,  and  the  carbon 
remains  behind  as  a  black  mass. 

The  same  fact  is  illustrated  by  adding  20  grams  of  cane  sugar 
to  15  cc.  of  hot  water  in  a  large  beaker.  Add  20  cc.  of  concen- 
trated sulphuric  acid.  A  vigorous  reaction  will  take  place,  steam 
escaping,  and  a  black  carbonaceous  mass  will  remain  behind. 

That  sulphuric  acid  is  a  strong  acid  can  be  shown  by  adding 
a  few  drops  of  the  dilute  acid  to  a  beaker  containing  100  cc.  of 
water  and  then  adding  a  few  drops  of  litmus.  The  litmus 
will  be  colored  deep  red. 

Perform  the  isame  experiment,  using  methyl  orange  as  the 
indicator.  What  do  you  observe  ?  Add  10  cc.  of  dilute  sul- 
phuric acid  to  each  of  two  beakers  containing  100  cc.  of  water. 
Neutralize  the  acid  in  one  beaker  with  sodium  hydroxide, 
and  the  acid  in  the  other  beaker  with  potassium  hydroxide, 
using  litmus  as  the  indicator.  Evaporate  the  two  solutions  to 
dry  ness.  What  salts  are  obtained  in  the  two  cases  ? 

Treat  dry  sodium  chloride  in  a  test-tube  with  concentrated 
sulphuric  acid,  hydrochloric  acid  is  evolved.  This  does  not  mean 
that  sulphuric  acid  is  stronger  than  hydrochloric  acid.  Indeed, 
sulphuric  acid  is  just  about  one-half  as  strong  as  hydrochloric 
acid,  as  has  been  shown  by  physical-chemical  methods  which 
it  would  lead  us  too  far  to  discuss  at  present.  Hydrochloric 
acid,  being  volatile,  is  set  free.  What  other  analogous  examples 
have  we  met  with  ?  (See  page  159.) 


PROBLEMS 

1.  With  what  weight  of  sulphur  would  10  litres  of  oxygen 
combine  to  form  sulphur  dioxide  ?  to  form  sulphur  trioxide  ? 

2.  Calculate  the  percentage  of  sulphur  in  ferrous  sulphide  ? 


SULPHUR,   SELENIUM,   TELLURIUM  171 

3.  With  what  weight  of  hydrogen  sulphide  would  5  litres  of 
oxygen  combine  to  form  sulphur  and  water? 

4.  How   much   zinc   sulphate,  lead  nitrate,  and   cadmium 
chloride  would  be  decomposed  by  1  gram  of  hydrogen  sulphide 
to  form  ZnS,  PbS,  and  CdS  ? 

5.  When  sulphuric  acid  acts  on  10  grams  of  copper,  what  is 
the  weight  of  the  sulphur  dioxide  set  free  ?    How  much  acid 
would  be  used  up  ? 

6.  How  much  sulphuric  acid  could  be  obtained  from  the 
sulphur  in  1.5  grams  of  pyrites  ? 

7.  Amount  of  ammonium  hydroxide  required  to  neutralize 
the  above  amount  of  sulphuric  acid  ? 


CHAPTER   XVI 

PHOSPHORUS,  ARSENIC,  ANTIMONY,  BISMUTH 
PHOSPHORUS  (At.  Wt.  =  31.0) 

Occurrence  and  Preparation.  —  Phosphorus  derives  its  name 
from  the  fact  that  it  emits  light,  or,  as  we  say,  is  phospho- 
rescent. It  does  not  occur  in  the  free  state,  but  mainly  in 
the  form  of  phosphates,  and  especially  in  combination  with 
calcium  as  the  calcium  salt.  This  is  the  compound  of 
phosphorus  which  occurs  with  other  substances  as  apatite, 
phosphorite,  etc.,  and  the  great  "  phosphate  beds  "  in  the  south- 
ern part  of  the  United  States  are  mainly  calcium  phosphate. 
Phosphorus  also  occurs  in  the  bones  of  animals  in  the  form 
of  the  calcium  salt.  Phosphorus  is  prepared  from  calcium 
phosphate.  If  bones  are  used  as  the  source  of  the  calcium 
phosphate,  the  organic  matter  is  first  destroyed  by  burning. 

The  tricalcium  phosphate  is  treated  with  sulphuric  acid, 
when  monocalcium  phosphate  is  formed :  — 

Ca3(P04)2  +  2  H2S04  =  2  CaS04  +  Ca(H2P04)2. 

This  is  then  heated,  when  it  passes  over  into  calcium  meta- 
phosphate :  — 

Ca(H2P04)2  =  2  H20  +  Ca(P03)2. 

The  calcium  metaphosphate  is  then  heated  with  a  mixture 
of  silicon  dioxide  (Si02)  and  powdered  charcoal.  The  follow- 
ing reaction  takes  place :  — 

Ca(P03)2  +  50  +  Si02  =  5  CO  +  CaSiO3  +  2  P. 

Properties  of  Phosphorus.  —  Phosphorus  is  a  soft  solid,  with 
a  slightly  yellowish  tint.  In  contact  with  the  air  it  combines 
readily  with  the  oxygen,  forming  an  oxide  of  phosphorus. 

172 


PHOSPHORUS,    ARSENIC,   ANTIMONY,    BISMUTH       173 

Phosphorus  combines  with  most  of  the  elements,  and  with 
such  elements  as  iodine  and  bromine  with  great  vigor. 

Phosphorus  is  an  extremely  poisonous  substance,  and  in 
working  with  it  precaution  must  be  taken  not  to  inhale  its 
vapors. 

Phosphorus  exists  in  more  than  one  form,  there  being  no  less 
than  four  allotropic  modifications.  Ordinary  yellow  phosphorus 
has  already  been  briefly  described. 

When  yellow  phosphorus  is  allowed  to  stand  under  water 
for  a  long  time,  exposed  to  the  light,  it  passes  over  into  a  red 
modification.  Red  phosphorus  is  easily  prepared  by  heating 
the  yellow  phosphorus  to  250°  in  an  atmosphere  free  from 
oxygen,  or  to  300°  in  a  vacuum  for  a  few  minutes.  Red 
phosphorus  is  an  amorphous  powder,  and  to  the  eye  resembles 
in  no  respect  the  ordinary  variety.  Red  phosphorus  is  much 
less  active  chemically  than  yellow.  When  heated  to  200°  in 
the  air,  it  does  not  take  fire.  When  brought  in  contact  with 
elements  and  compounds  with  which  yellow  phosphorus  unites 
at  once,  it  does  not  combine  with  them. 

When  red  phosphorus  is  heated  to  260°  in  an  atmosphere  of 
carbon  dioxide,  it  passes  over  quantitatively  into  the  yellow 
modification. 

We  have  in  these  two  varieties  of  phosphorus  a  case 
somewhat  analogous  to  that  met  with  in  the  two  modifica- 
tions of  oxygen  and  of  sulphur.  We  saw  in  the  cases  of 
oxygen  and  sulphur  that  the  real  difference  between  the 
properties  of  the  two  modifications  was  to  be  sought  for  in 
the  different  amounts  of  intrinsic  energy  present  in  the  two 
modifications. 

We  should,  therefore,  naturally  ask  whether  there  is  any 
similar  relation  between  the  two  modifications  of  phosphorus 
described  above  ?  Do  the  different  modifications  contain  dif- 
ferent amounts  of  intrinsic  energy  ?  This  can  be  answered  by 
burning  the  different  modifications  in  oxygen,  when  they  yield 
the  same  end  product,  phosphorus  pentoxide,  P205.  The 
results  of  thermochemical  measurements  show  that  yellow 


174  ELEMENTS   OF   INORGANIC   CHEMISTRY 

phosphorus  when  burned  evolves,  per  gram-atomic  weight, 
27,300  calories  of  heat  more  than  red  phosphorus,  and  this  is 
approximately  the  thermochemical  equivalent  of  the  difference 
between  the  intrinsic  energies  of  these  two  modifications  of 
phosphorus. 

Hydrogen  Phosphide,  or  Phosphine,  is  prepared  by  the  action 
of  caustic  potash  on  phosphorus  in  the  presence  of  water :  — 

3  KOH  +  4  P  +  3  H20  =  3  H2KP02  +  PH3. 

Phosphine  produced  by  this  method  always  contains  a  little 
of  the  liquid  compound  P2H4,  which  renders  it  spontaneously 
inflammable. 

Compounds  of  Phosphorus  with  Oxygen  and  Hydrogen. — 
Phosphorus  forms  a  number  of  compounds  with  oxygen. 
The  more  important  are:  phosphorus  sesquioxide,  P203,  and 
phosphorus  pentoxide,  P205. 

The  sesquioxide  of  phosphorus,  P203,  is  formed  by  the 
incomplete  oxidation  of  phosphorus.  When  phosphorus  is 
burned  in  a  slow  current  of  air,  which  does  not  furnish  enough 
oxygen  to  convert  it  into  the  pentoxide,  it  forms  the  ses- 
quioxide. The  sesquioxide  readily  takes  up  oxygen  and  passes 
over  into  the  pentoxide. 

Phosphorus  pentoxide,  P205,  is  formed  by  the  oxidation  of 
phosphorus  in  the  presence  of  an  excess  of  oxygen.  It  is  a 
beautifully  white  compound,  which  has  remarkable  power  to 
combine  with  water.  Indeed,  it  is  the  best  drying  agent  at  the 
disposal  of  the  chemist.  Phosphorus  pentoxide  is  the  anhy- 
dride of  an  acid.  When  it  combines  with  the  maximum 
amount  of  water,  it  forms  phosphoric  acid  :  — 

PA  +  3H20  =  2H3P04. 

This  brings  us  to  the  acids  of  phosphorus,  of  which  there 
are  several. 

The  Acids  of  Phosphorus.  —  Phosphorus  combines  with 
oxygen  and  hydrogen,  forming  a  number  of  compounds  which 
are  acids.  The  more  important  are  ;  — 


PHOSPHORUS,    ARSENIC,   ANTIMONY,   BISMUTH       175 

Phosphoric  acid        ....  H3P04 

Pyrophosphoric  acid        .         .         .  H4P2O7 

Metaphosphoric  acid        .         .         .  HP03 

Phosphorous  acid     ....  H3P03 

The  most  important  of  these,  by  far,  is  ordinary  phosphoric 
acid,  or  orthophosphoric  acid. 

Orthophosphoric  Acid,  H3P04.  —  Orthophosphoric  acid  is 
formed,  as  already  stated,  by  dissolving  phosphorus  pentoxide 
in  water.  It  is  in  the  form  of  salts  of  this  acid  that  phos- 
phorus occurs  in  nature ;  the  calcium  salt,  Ca3P04,  being  the 
compound  in  which  phosphorus  occurs  in  the  great  phosphate 
beds.  When  this  salt,  which  is  insoluble  in  water,  is  treated 
with  an  excess  of  concentrated  sulphuric  acid,  the  following 
reaction  takes  place :  — 

Ca3(P04)2  +  3  H2S04  =  3  CaS04  +  2  HgPO,, 

giving  free  orthophosphoric  acid. 

The  tertiary  or  normal  phosphates  have  the  composition 
M3P04,  where  M  is  a  univalent  metal;  the  secondary  phos- 
phates, the  composition  HM2P04 ;  and  the  primary  phosphates, 
the  composition  H2MP04. 

Dissociation  of  Phosphoric  Acid.  —  Since  phosphoric  acid  is 
a  tribasic  acid,  it  must  dissociate  into  three  hydrogen  ions. 
The  complete  dissociation  of  phosphoric  acid  in  the  following 
sense  is  very  difficult  to  effect :  — 

H3P04  =  H,  H,  H,  P04, 

since  phosphoric  acid  is  a  comparatively  weak  acid. 
Phosphoric  acid  dissociates  first  as  follows :  — 

H3P04  =  H,H2P04. 

When  more  water  is  added,  or  when  these  hydrogen  ions 
have  been  used  up  by  a  base,  the  second  hydrogen  begins  to 
split  off  in  the  ionic  state  :  — 

H2P04  =  H,  HP04. 


176  ELEMENTS  OF   INORGANIC   CHEMISTRY 

It  is  not  until  these  hydrogen  ions  have  been  used  up,  or  very 
great  dilution  has  been  reached,  that  the  third  hydrogen  ions 
begin  to  split  off :  —  =  + 

HP04  =  H,  P04. 

Hydrolytic  Dissociation.  —  It  is  a  general  rule  that  the  salts 
of  weak  acids  or  bases  are  acted  on  by  water  to  a  greater  or 
less  extent,  being  broken  down  into  the  corresponding  acid 
and  base.  Take  tertiary,  or  normal  sodium  phosphate, 
Na3P04.  When  this  is  acted  on  by  water,  the  following  decom- 
position takes  place  to  some  extent :  — 

Na^O,  +  H20  =  Na,   Na,   HP04  +  Na,   OH. 

This  kind  of  dissociation  is  known  as  hydrolytic  dissociation. 
This  is  the  explanation  of  the  alkaline  reaction  shown  by  such 
compounds  in  water. 

Pyrophosphoric  Acid,  H4P207.  —  Pyrophosphoric  acid  is 
formed  from  phosphoric  acid  by  loss  of  water :  — 

2H3P04  =  H20  4-  H4P207. 

Salts  of  this  acid  are  easily  obtained  by  heating  secondary 

phosphates :  — 

2  HM2P04  =  H,0  +  M4P207. 

Metaphosphoric  Acid,  HP03.  — Metaphosphoric  acid  is  formed 
when  normal  phosphoric  acid  is  heated  higher  than  is  neces- 
sary to  form  the  pyroacid :  — 

H3P04  =  H20  +  HP03. 

Salts   of   this    acid    are   formed  when   primary  phosphates, 
MH2P04,  are  heated  :  — 

MH2PO4  =  H20  +  MP03. 

Phosphorous  Acid,  H3P03.  —  Phosphorous  acid  is  formed  by 
the  action  of  phosphorus  on  moist  air.  Also  by  the  action  of 
water  on  a  chloride  of  phosphorus  with  which  we  shall  soon 
become  familiar  —  phosphorus  trichloride:  — 

PC13  +  3  H20  =  3  HC1  +  H3P03. 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   BISMUTH       177 

Phosphorous  acid  contains  three  hydrogen  atoms,  and  would, 
therefore,  be  expected  to  be  a  tribasic  acid.  The  fact  is,  it  is 
only  dibasic,  the  salts  richest  in  metal  having  the  composition 
M2HP03.  This  is  to  be  explained  as  follows :  The  first  stage 
in  the  dissociation  of  this  substance  is  represented  by  the 
following  equation :  — 

H3P03  =  H,  H2P03. 

The  second  stage  is  represented  thus :  — 

H2P03  =  H,  HP03. 

It  is  impossible  to  go  farther  and  split  off  the  last  hydrogen 
atom  as  an  ion. 


COMPOUNDS   OF  PHOSPHORUS   WITH  THE  HALOGENS 

Phosphorus  Trichloride,  PC13.  —  When  chlorine  gas  is  passed 
over  an  excess  of  phosphorus,  in  an  atmosphere  free  from 
oxygen,  the  two  combine  and  form  phosphorus  trichloride :  — 

2P  +  3C12  =  2PC13. 

This  compound,  when  brought  in  contact  with  water,  decom- 
poses, forming  phosphorous  acid  and  hydrochloric  acid :  — 

PC13  +  3  H20  =  3  HC1  +  H3P03. 

Phosphorus  has  the  power  to  take  up  more  chlorine  and  form 
phosphorus  pentachloride. 

Phosphorus  Pentachloride,  PC15.  —  The  pentachloride  of  phos- 
phorus is  formed,  as  stated  above,  by  the  action  of  chlorine  on 
the  trichloride  of  phosphorus ;  also  by  the  direct  action  of  an 
excess  of  chlorine  on  phosphorus.  Like  the  trichloride  it  is 
readily  decomposed  by  water,  forming  phosphoric  and  hydro- 
chloric acids :  — 

PC15  +  4  H20  =  5  HC1  +  H3P04. 


178  ELEMENTS   OF   INORGANIC   CHEMISTRY 

ARSENIC  (At.   Wt.  =  74.9) 

Occurrence,  Preparation,  and  Properties.  —  An  element  closely 
allied  chemically  to  phosphorus  is  arsenic.  That  there  are 
marked  differences,  however,  will  appear  as  the  following 
chapter  develops. 

Arsenic  does  occur  in  the  free  condition.  It  is  generally  in 
combination  with  the  metals,  either  directly  as  in  the  compound 
with  iron,  Fe2As3,  or  with  sulphur  as  in  arsenical  pyrites,  FeAsS. 

Arsenic  is  generally  obtained  from  its  compounds  by  simply 
heating  them ;  the  arsenic  being  volatile,  passes  off  as  vapor. 

Arsenic  is  a  solid  at  ordinary  temperatures,  gray  in  color  and 
very  brittle. 

Compound  of  Arsenic  with  Hydrogen  —  Arsine,  AsH  — 
Arsenic  forms  with  hydrogen  the  compound  AsH3,  which  is 
analogous  to  ammonia,  NH3.  It  is  formed  by  the  action  of 
nascent  hydrogen  on  compounds  of  arsenic.  When  we  have  a 
•compound  of  arsenic  in  the  presence  of  zinc  and  an  acid,  arsine 
is  formed.  If  the  compound  is  ordinary  arsenic  trioxide,  As203, 
this  is  reduced  by  nascent  hydrogen  as  follows :  — 

As203  +  6  H2  =  3  H20  +  2  AsH3. 

When  arsine  is  heated,  it  is  broken  down  into  its  elements, 
arsenic  and  hydrogen.  When  arsine  is  burned  and  a  cold  object 
introduced  into  the  flame,  arsenic  is  deposited  upon  the  object. 
These  reactions  are  made  use  of  for  the  detection  of  arsenic. 
Marsh's  method  for  detecting  arsenic  consists  in  reducing  the 
arsenic  compound  to  arsine  and  burning  the  arsine. 


COMPOUNDS  OF   ARSENIC   WITH   OXYGEN  AND 
HYDROGEN 

Compounds  of  Arsenic  with  Oxygen. — Arsenic  forms  two  com- 
pounds with  oxygen.  One  of  these  has  the  composition  As-jO^ 
and  is  known  as  arsenic  trioxide;  the  other  has  the  compo- 
sition As205,  and  is  known  as  arsenic  pentoxide. 


PHOSPHORUS,   ARSENIC,    ANTIMONY,   BISMUTH       179 

Arsenic  Trioxide,  As203,  is  formed  when  arsenic  is  oxidized 
either  by  burning  in  the  air,  or  by  some  strong  oxidizing  agent 
such  as  nitric  acid. 

Arsenic  Pentoxide,  As205.  —  Arsenic  pentoxide  cannot  be 
formed  like  phosphorus  pentoxide  by  burning  the  element  in 
oxygen.  It  is  prepared  by  removing  water  from  arsenic  acid, 
and  is,  therefore,  the  anhydride  of  this  acid. 

Arsenious  Acid,  H3As03.  —  This  acid  is  not  known  in  the  free 
condition.  There  are,  however,  three  classes  of  salts  known, 
depending  upon  whether  one,  two,  or  three  of  the  hydrogen 
ions  have  given  their  electrical  charges  to  the  metal  atoms 
which  have  entered  the  compound  in  their  stead. 

Arsenic  Acid,  H3As04.  —  When  ordinary  white  arsenic,  or 
arsenic  trioxide,  is  heated  with  some  strong  oxidizing  agent 
such  as  nitric  acid,  or  aqua  regia,  in  the  presence  of  water,  it  is 
oxidized  to  arsenic  acid.  The  reaction  consists  in  the  direct 
addition  of  oxygen  and  water :  — 

As203  +  02  +  3  H20  =  2  H3  As04. 

The  acid  is  known  in  solution  as  a  syrupy  liquid,  and  in  the 
solid  form  as  white  needles.  Like  phosphoric  acid  it  forms 
three  series  of  salts, — primary,  secondary,  and  tertiary. 

Sulpho-salts  of  Arsenic.  —  Arsenic  forms  with  sulphur  and  the 
alkali  metals,  salts  of  acids  having  the  composition  H3AsS3  and 
H3AsS4.  These  acids  are  the  sulphur  analogues  of  arsenious 
acid,  H3As03,  and  arsenic  acid,  H3As04.  The  sulpho-acids  or 
thioacids  are  themselves  not  known,  but  certain  salts  are  well- 
characterized  substances. 

ANTIMONY  (At.  Wt.  120.0) 

Occurrence  and  Preparation.  —  Another  element  which  pre- 
sents many  chemical  analogies  to  phosphorus  and  arsenic  is 
antimony.  Antimony  occurs  in  nature  chiefly  as  the  trisul- 
phide,  Sb2S3,  which  is  known  as  the  mineral  stibnite.  It  also 
occurs  in  combination  with  arsenic  and  also  with  oxygen. 


180  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Antimony  is  prepared  from  stibnite  by  roasting  out  the  sul- 
phur. The  sulphide  is  heated  in  the  air,  when  the  sulphur  is 
converted  into  the  dioxide,  and  the  antimony  into  the  trioxide 
or  sesquioxide,  Sb203. 

The  oxide  is  then  reduced  with  carbon  :  — 


3C  =  3C02  +  4Sb. 

Compound  of  Antimony  with  Hydrogen  —  Stibine,  SbH3.  — 

Antimony  forms  with  hydrogen  the  compound  SbH3,  which  is 
analogous  to  the  compounds  of  hydrogen  with  nitrogen,  phos- 
phorus, and  arsenic. 

NH3  ......  ammonia 

PH3  ......  phosphine 

AsH3  ......  arsine 

SbH3  .         .         .  .         .  stibine 

COMPOUNDS   OF   ANTIMONY   WITH    OXYGEN  AND 
HYDROGEN 

With  oxygen  alone  antimony  forms  three  compounds:  Anti- 
mony sesquioxide  —  Sb203,  antimony  tetroxide  —  Sb204,  and 
antimony  pentoxide  —  Sb205. 

Acids  of  Antimony.  —  Antimony  combines  with  hydrogen  and 
oxygen,  forming  several  compounds,  which  are  acids,  but  these 
are  not  so  numerous  as  in  the  cases  of  phosphorus  and  arsenic. 
The  best  known  of  these  is  antimonic  acid,  having  the  compo- 
sition H3Sb04.  This  is  formed  by  the  action  of  strong  oxidizing 
agents  such  as  nitric  acid,  or  aqua  regia,  on  antimony.  It  is 
also  formed  when  the  pentachloride  of  antimony  is  treated  with 
water,  which  is  analogous  to  the  formation  of  phosphoric  acid 
from  phosphorus  pentachloride  :  — 

SbCl5  +  4  H20  =  5  HC1  +  H3Sb04. 

Compounds  of  Antimony  with  Sulphur  and  the  Metals.  —  We 

have  seen  that  arsenic  combines  with  sulphur  and  the  metals, 
forming  salts  of  sulpho-acids  of  arsenic.  In  an  analogous 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   BISMUTH       181 

manner,  antimony  forms  salts  of  sulpho-acids.  When  anti- 
mony trisulpjiide  is  treated  with  the  sulphide  of  an  alkali  metal, 
such  as  potassium  sulphide,  ammonium  sulphide,  or  poly  sul- 
phide, the  antimony  trisulphide  dissolves,  forming  a  salt 
(MSbS2)  of  meta-sulphantimonious  acid  (HSbS2).  The  salt  is 
a  meta-sulphantimonite.  Salts  of  sulphantimonious  acid  (H3SbS3) 
are  also  known. 

When  antimony  pentasulphide  is  dissolved  in  an  alkaline 
sulphide,  salts  of  sulphantimonic  acid  are  formed :  — 

Sb2S5  +  3  Na2S  =  2  Na3SbS4. 

This  compound,  which  contains  nine  molecules  of  water,  is 
known  as  Schlippe's  salt. 


BISMUTH  (At.  Wt.  =  208.3) 

The  last  member  of  the  nitrogen,  phosphorus,  arsenic,  anti- 
mony family  of  group  V  in  the  Periodic  System  is  bismuth. 
We  have  seen  that  as  the  atomic  weight  increases,  the  elements 
become  less  acidic,  and  the  basic  properties  begin  to  manifest 
themselves.  This  condition,  which  has  already  appeared  in 
antimony,  is  intensified  in  the  element  which  we  are  now  about 
to  study,  —  bismuth. 

Occurrence  and  Properties. — Bismuth  occurs  mainly  in  the 
free  condition,  but  also  combined  with  sulphur  as  the  trisul- 
phide, Bi2S3.  Bismuth  is  obtained  from  the  sulphide  by  burn- 
ing out  the  sulphur  with  oxygen,  when  it  is  transformed  into 
the  oxide.  The  oxide  is  then  reduced  by  carbon,  yielding  the 
element. 

Some  of  the  most  important  substances  containing  bismuth 
are  certain  of  its  alloys  with  other  metals.  These  have  the 
remarkable  property  that  they  fuse  far  below  the  melting-point 
of  the  lowest-melting  constituent.  The  well-known  Rose's 
fusible  metal  consists  of  one  part  of  lead,  one  part  of  tin,  and 
two  parts  of  bismuth.  It  fuses  at  93°.8. 

There  is  an  alloy  of  bismuth  still  more  remarkable  than  the 


182  ELEMENTS   OF   INORGANIC   CHEMISTRY 

above,  in  that  it  fuses  at  60°.5.  It  is  known  as  Wood's  metal. 
This  contains  two  parts  of  lead,  one  part  of  tin,  four  parts  of 
bismuth,  and  one  part  of  cadmium.  It  is  the  lowest-melting 
of  these  substances. 

Compounds  of  Bismuth  with  Oxygen  and  Hydrogen. — While 
bismuth  forms  four  compounds  with  oxygen,  the  only  com- 
pound of  importance  is  the  sesquioxide,  Bi^.  It  is  formed 
when  bismuth  burns  in  the  air.  It  combines  readily  with  acids, 
forming  water  and  the  corresponding  salt,  and  is,  therefore,  a 
base. 

The  corresponding  hydroxide,  Bi(OH)3,  also  has  decidedly 
basic  properties. 

Certain  rare  elements  in  this  group  of  the  Periodic  System 
are:  vanadium,  columbium,  neodymium,  praseodymium,  and 
tantalum. 

EXPERIMENTS  WITH  PHOSPHORUS,  ARSENIC, 
ANTIMONY,  AND  BISMUTH 

Experiment  100.    Phosphorus  an  Active  Element  Chemically. 

—  (Porcelain  crucible ;  stone  slab;  filter-paper;  carbon  bisul- 
phide ;  iodine ;  phosphorus.) 

In  a  porcelain  crucible  resting  on  a  stone  slab  under  a  hood 
introduce  a  piece  of  phosphorus  dried  by  pressing  between 
filter-paper.  (Never  touch  phosphorus  with  the  hand,  or  severe 
burns  may  result.)  Allow  the  phosphorus  to  stand  exposed  to 
the  air.  White  fumes  will  soon  escape  from  the  phosphorus, 
and  its  temperature  may  become  high  enough  to  cause  it  to 
ignite. 

Finely  divided  phosphorus  always  takes  fire  in  contact  with 
the  air.  Dissolve  a  piece  of  phosphorus  the  size  of  a  pea  in 
2  cc.  of  carbon  bisulphide.  Seize  a  piece  of  filter-paper  with  a 
pair  of  forceps  and  introduce  it  into  the  solution.  Lay  the 
filter-paper  on  a  stone  slab  to  dry.  The  carbon  bisulphide 
quickly  evaporates,  leaving  the  phosphorus  on  the  paper  in  a 
very  finely  divided  condition.  It  always  takes  fire  spontane- 
ously under  these  conditions. 

That  phosphorus  combines  ivith  iodine  by  mere  contact  can 
be  seen  by  bringing  into  the  porcelain  crucible  a  piece  of  dry 
phosphorus  the  size  of  a  pea,  and  throwing  on  to  the  phos- 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   BISMUTH       183 

phorus  a  few  crystals  of  iodine.  The  two  combine  at  once  on 
contact.  (See  page  173.) 

Experiment  101.  Red  Phosphorus  more  Stable  than  Yellow. 
—  (Porcelain  crucible;  strip  of  metal  25  cm.  long;  iodine;  red 
and  yellow  phosphorus.) 

Bring  a  little  iodine  in  contact  with  a  little  red  phosphorus. 
They  do  not  react.  On  the  middle  of  a  strip  of  metal,  25  cm. 
long,  held  at  one  end  by  a  clamp,  place  side  by  side  a  piece  of 
yellow  and  a  piece  of  red  phosphorus.  Heat  the  other  end 
of  the  strip.  The  yellow  phosphorus  will  take  fire  first. 
Repeat  the  experiment,  placing  the  red  phosphorus  10  cm.  from 
the  end  to  be  heated,  and  the  yellow  phosphorus  20  cm.  from 
the  end.  Even  under  these  conditions  the  yellow  phosphorus 
will  take  fire  first.  (See  page  173.) 

Experiment  102.  Compounds  of  Phosphorus  and  Hydrogen, 
Preparation.  —  (Glass  flask  holding  150  cc. ;  glass  tubing ; 
crystallizing-dish  ;  sodium  hydroxide ;  yellow  phosphorus ; 
hydrogen  generator.) 


FIG.  48. 


Set  up  the  apparatus  as  in  Figure  48  (under  the  hood). 
Ten  grams  of  caustic  potash  are  dissolved  in  25  cc.  of  water, 
and  the  mixture  cooled  in  ice-water.  It  is  then  introduced  into 
the  flask  and  two  or  three  small  fragments  of  phosphorus  added. 
The  flask  is  then  closed  tightly  with  the  stopper.  The  tube 


184  ELEMENTS  OF  INORGANIC  CHEMISTRY 

D  is  then  attached  to  the  hydrogen  generator  (a  Kipp 
apparatus  if  convenient),  and  hydrogen  is  swept  through  the 
apparatus  until  all  of  the  air  has  been  removed.  The  rubber 
tube  placed  on  the  end  of  D  is  closed  with  a  pinch-cock,  and  the 
hydrogen  generator  removed.  The  contents  of  the  flask  are 
then  gently  heated.  The  hydrogen  is  first  driven  out  of  the 
flask;  then  phosphine  begins  to  escape,  and  as  each  bubble 
comes  in  contact  with  the  air  it  takes  fire  spontaneously,  pro- 
ducing a  ring  of  white  fumes  which  rises  in  the  air.  If  the 
experiment  is  carried  out  in  a  place  free  from  strong  air-cur- 
rents, the  effect  is  very  beautiful.  The  white  fumes  rise  in  a 
continuous  ring,  showing  distinct  vortex  movement. 

After  the  experiment  is  ended  remove  the  flame  and  allow 
the  water  to  flow  back  from  the  dish  into  the  flask.  Any  gas 
remaining  in  the  flask  should  be  driven  out  by  hydrogen  before 
the  flask  is  opened. 

The  inflammability  is  not  due  to  the  compound  PH3.  If  the 
gases  escaping  from  the  flask  are  washed  by  passing  through  a 
cylinder  containing  alcohol,  they  are  no  longer  spontaneously 
inflammable  on  coming  in  contact  with  the  air.  The  alcohol 
dissolves  the  compound  P2H4,  which  is  also  formed  in  the 
above  experiment,  and  which  is  the  spontaneously  inflammable 
constituent. 

If  in  the  above  experiment  considerable  alcohol  is  added  to 
the  flask,  this  dissolves  the  compound  P2H4  as  fast  as  it  is 
formed,  and  the  escaping  gas  is  no  longer  inflammable.  (See 
page  174.) 

Experiment  103.  Phosphoric  Acid  formed  when  Water  acts 
on  Phosphorus  Pentoxide.  —  (Phosphorus  pentoxide;  water.) 

Add  100  cc.  of  distilled  water  to  a  beaker,  and  introduce,  by 
means  of  a  small  porcelain  spatula,  quantities  of  phosphorus 
pentoxide  not  larger  than  a  pea.  Continue  to  add  the  pentoxide 
until  a  gram  or  so  has  dissolved  in  the  water. 

Filter  a  part  of  the  solution ;  this  contains  chiefly  metaphos- 
phoric  acid.  Test  the  reaction  of  the  solution  towards  litmus. 

Boil  the  remainder  of  the  solution,  and  the  undissolved  por- 
tion will  pass  into  solution.  This  is  a  solution  of  orthophos- 
phoric  acid.  Test  its  reaction  with  litmus.  (See  page  175.) 

Experiment  104.  Pyrophosphate  formed  by  heating  a 
Secondary  Phosphate.  —  (Porcelain  crucible ;  secondary  sodium 
phosphate ;  porcelain  crucible.) 

Introduce  a  few  grams  of  secondary  sodium  phosphate  into 
a  porcelain  crucible,  and  heat  gently  over  a  bunsen  burner  until 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   BISMUTH       185 

all  the  water  is  removed.  Then  heat  sharply  over  the  blast- 
lamp.  Dissolve  the  pyrophosphate  in  water  and  add  a  solution 
of  silver  nitrate.  What  do  you  observe  ?  (See  page  176.) 

Experiment  105.  Metaphosphate  formed  by  heating  Micro- 
cosmic  Salt  (Sodium  Ammonium  Acid  Phosphate,  NaNH4, 
HP04).  —  Heat  a  few  grams  of  microcosmic  salt  in  a  porcelain 
crucible,  first  over  a  bunsen  burner  and  then  over  a  blast-lamp. 
Sodium  metaphosphate  results.  (Equation?)  Dissolve  in  water 
and  add  silver  nitrate.  What  do  you  observe  ?  (See  page  176.) 

Experiment  106.  Arsenic  sublimes  without  melting;  burns 
on  the  Air;  formed  by  reducing  the  Oxide  with  Carbon. — 
(Hard-glass  tube,  15  cm.  long,  closed  at  one  end ;  arsenic.) 

Introduce  one  or  two  small  pieces  of  arsenic  into  a  hard-glass 
tube  closed  at  one  end,  and  heat  this  portion  of  the  tube  con- 
taining the  arsenic.  The  arsenic  will  vaporize  over  the  flame 
without  melting,  but  will  condense  again  on  the  colder  walls  of 
the  tube. 

Heat  a  small  piece  of  arsenic  on  a  piece  of  charcoal  over  the 
blast-lamp.  White  fumes  of  arsenic  trioxide  will  be  formed  in 
abundance. 

All  of  the  experiments  with  arsenic  must  be  carried  out  under 
a  hood  with  a  good  draft.  TJie  fames  of  arsenic  and  its  com- 
pounds are  deadly  poisonous  if  inhaled. 

Mix  a  little  arsenic  trioxide  —  white  arsenic  —  with  an  equal 
volume  of  powdered  charcoal,  and  introduce  the  mixture  into  a 
hard-glass  tube  closed  at  one  end,  through  a  funnel  to  prevent 
any  of  the  mixture  from  adhering  to  the  walls  of  the  tube. 

Heat  the  mixture ;  arsenic  will  be  set  free,  and  will  distil  out 
of  this  portion  of  the  tube  and  condense  on  the  colder  walls  of 
the  tube.  (See  page  178.) 

Experiment  107.  Preparation  of  Arsine.  —  (Flask  holding 
500  cc. ;  safety-tube ;  calcium  chloride  tube ;  hard-glass  tube 
drawn  out  as  in  Figure  48 ;  pure  zinc ;  dilute  hydrochloric  acid ; 
white  arsenic.) 

The  apparatus  to  be  used  is  shown  in  Figure  49.  The  flask 
A  contains  pure  zinc  to  which  dilute  hydrochloric  acid  is  added. 
The  hydrogen  generated  is  passed  through  the  tube  filled  with 
calcium  chloride  to  dry  it.  After  all  the  air  has  been  driven 
out  of  the  apparatus  the  hydrogen  is  ignited  at  B.  It  burns 
with  a  very  nearly  colorless  flame,  since  the  tube  B  is  of  hard 
glass.  Introduce  a  piece  of  cold  porcelain  into  the  hydrogen 
flame  ;  it  will  not  be  colored. 

Now  introduce  into  the  flask  A,  through  the  safety-tube,  a 


186 


ELEMENTS   OF  INORGANIC   CHEMISTRY 


few  drops  of  a  solution  of  arsenic  trioxide  in  hydrochloric  acid. 
The  appearance  of  the  flame  of  burning  hydrogen  will  be 
changed  at  once.  It  will  acquire  a  milky-white  hue,  and  white 
fumes  will  escape.  Introduce  again  the  cold  porcelain,  and  a 


FIG.  49. 


dark  spot  will  be  deposited  upon  it.  This  is  the  so-called 
arsenic  mirror.  Observe  carefully  its  metallic  appearance. 
Pour  over  the  arsenic  mirror  a  solution  of  sodium  hypochlorite. 
The  arsenic  will  be  dissolved.  (See  page  178.) 

Experiment  108.  Arsine  decomposed  by  Heat.  —  Heat  one 
of  the  expansions  in  the  glass  tube  (Fig.  48).  The  arsine  will 
be  decomposed,  and  a  mirror  of  arsenic  will  be  deposited  upon 
the  glass  just  beyond  the  flame.  Interrupt  the  experiment, 
open  the  glass  tube,  and  add  a  solution  of  sodium  hypochlorite 
and  shake.  Does  the  arsenic  dissolve  ?  (See  page  178.) 

Experiment  109.  Arsenic  Trisulphide.  —  (Test-tube ;  hydro- 
gen sulphide  generator ;  white  arsenic ;  hydrochloric  acid ; 
yellow  ammonium  sulphide.) 

Dissolve  a  little  white  arsenic  in  hydrochloric  acid  and  pass 
a  stream  of  hydrogen  sulphide  through  the  solution.  Yellow 
arsenious  sulphide  is  precipitated.  Write  the  equation  ? 

Filter  off  the  solid  sulphide  and  wash  with  hot  water.    To  the 


PHOSPHORUS,   ARSENIC,   ANTIMONY,   BISMUTH       187 

solid  sulphide  on  the  filter-paper  add  yellow  ammonium  sul- 
phide. The  arsenic  sulphide  passes  into  solution,  forming  the 
sulpho  salt.  Write  the  equation  ?  (See  page  179.) 

Experiment  110.  Fusion  of  Antimony  ;  Combustion  of  Anti- 
mony in  Contact  with  the  Air.  —  (Antimony  j  piece  of  charcoal ; 
blowpipe.) 

When  a  piece  of  antimony  is  heated  on  charcoal  by 
means  of  a  mouth-blowpipe  it  melts.  When  the  molten 
globule  is  thrown  on  a  piece  of  white  paper,  it  runs  over  the 
surface  of  the  paper,  and  often  produces  beautiful,  fantastic 
designs.  If  the  piece  of  charcoal  on  which  the  antimony  was 
melted  be  examined,  it  will  be  found  to  contain  a  white  coat- 
ing which  is  an  oxide  of  antimony.  Compare  this  with  the 
coating  produced  by  arsenic  ?  It  is  less  volatile  than  the  oxide 
of  arsenic,  and  remains  on  the  charcoal  closer  to  the  flame  of 
the  blowpipe.  (See  page  180.) 

Experiment  111.  Preparation  and  Properties  of  Stibine. — 
(Same  materials  as  in  Experiment  107 ;  instead  of  arsenic 
trioxide  use  a  solution  of  tartar  emetic  or  of  antimonious 
chloride,  to  which  some  tartaric  acid  has  been  added.) 

Arrange  the  apparatus  exactly  as  in  Figure  48,  and  prepare 
stibine  exactly  as  arsine  was  prepared.  Burn  the  mixture  of 
hydrogen  and  stibine  and  note  the  color  of  the  flame.  Introduce 
a  cold  piece  of  porcelain,  and  a  dark  spot  will  be  deposited  — 
the  antimony  mirror.  The  spot  of  antimony  is  dull  in 
appearance,  and  thus  differs  from  the  bright  lustrous  spot  of 
arsenic.  The  antimony  mirror  is  insoluble  in  sodium  hypo- 
chlorite,  and  this  distinguishes  it  unquestionably  from  arsenic. 

When  the  tube  through  which  the  stibine  is  passing  is 
heated,  the  stibine  is  decomposed  and  antimony  is  deposited 
on  the  walls  of  the  tube.  The  antimony  being  less  volatile 
than  the  arsenic  is  deposited  nearer  the  flame.  (See  page  180.) 

Experiment  112.  Antimony  Sulphide.  —  (Solution  of  tartar 
emetic ;  hydrogen  sulphide  generator.) 

Pass  hydrogen  sulphide  into  a  solution  of  tartar  emetic  to 
which  a  few  drops  of  hydrochloric  acid  have  been  added. 
Brick-red  antimony  sulphide  is  precipitated.  Filter  the  solid 
from  the  solution  and  wash  the  precipitate  on  the  filter-paper. 
Add  to  the  filter-paper  containing  the  antimony  sulphide  some 
yellow  sulphide  of  ammonium.  The  sulphide  of  antimony  will 
dissolve.  What  is  formed  ?  Write  equation?  (See  page  181.) 

Experiment  113.  Oxidation  of  Bismuth  on  the  Air.  —  (Piece 
of  charcoal ;  blowpipe ;  bismuth.) 


188  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Heat  a  piece  of  bismuth  on  charcoal-  with  the  blowpipe. 
What  is  the  color  of  the  oxide  formed  on  the  charcoal  ?  Is  it 
very  volatile  ?  Compare  with  the  results  obtained  with  arsenic 
and  antimony. 

Experiment  114.  Bismuth  Sulphide.  —  (Bismuth  nitrate; 
nitric  acid;  hydrogen  sulphide  generator.) 

Dissolve  bismuth  nitrate  in  water  containing  nitric  acid  and 
pass  hydrogen  sulphide  into  the  solution.  Black  bismuth  sul- 
phide is  precipitated.  Filter  off ;  test  solubility  in  dilute  hydro- 
chloric and  nitric  acids  ?  In  concentrated  hydrochloric  and 
nitric  acids  ? 


PROBLEMS 

1.  How  much  phosphorus  could  be  obtained  from  a  kilo- 
gram of  bone-ash,  if  bone-ash  were  pure  tricalcium  phosphate  ? 
Amount  of  phosphine  obtainable  from  the  above  amount  of 
phosphorus  ? 

2.  How  much  phosphoric  acid  can  be  obtained  from  5  kilo- 
grams of  bone-ash,  assuming  that  bone-ash  is  pure  tricalcium 
phosphate  ? 

3.  Weight  of  hydrogen  ions  obtainable  from  10  grams  of 
orthophosphoric  acid,  if  the  dissociation  of  each  molecule  into 
three  hydrogen  ions  was  complete  ? 

4.  When   3.7   grams    of    secondary   sodium    phosphate    is 
heated,  how   much  pyrophosphate  would  be  formed?     How 
much  metaphosphate  would  be  formed  from  the  same  weight 
of  primary  sodium  phosphate  ? 

5.  Weight  of  arsenic  trioxide  that  can  yield  7.9  grams  of 
arsine  ? 

6.  Weight  of  stibine  obtainable  from  the  antimony  in  3.1 
grams  of  stibnite  ? 

7.  The  bismuth  in  5  grams  of  the  compound  Bi2S3  will  form 
what  weight  of  the  hydroxide  Bi(OH)3  ? 


CHAPTER   XVII 

CARBON,    SILICON,   BORON 
CARBON  (At.  Wt.  =  12.0) 

WE  now  come  to  one  of  the  most  important  elements  in  the 
whole  field  of  chemistry.  This  is  the  first  member  of  group 
IV — Carbon.  This  element  is  important  as  being  an  essential 
constituent  of  every  living  thing,  from  the  simplest  organism 
to  the  most  complex. 

AUotropic  Forms  of  Carbon:  Diamond  and  Graphite.  —  We 
know  carbon  in  several  modifications,  both  crystallized  and 
amorphous.  There  are  two  crystalline  modifications,  known 
respectively  as  diamond  and  graphite.  The  diamond  is  carbon 
and  nothing  but  carbon,  as  is  shown  by  the  fact  that  when  the 
diamond  is  burned  in  oxygen,  it  is  converted  completely  into 
a  compound  of  carbon ;  and  when  this  compound  is  collected 
and  weighed,  the  amount  of  carbon  present  in  the  compound  is 
exactly  equal  to  the  weight  of  the  original  diamond. 

Diamonds  have  now  been  made  artificially  —  a  problem 
which  has  attracted  great  attention  in  time  past.  The  French 
chemist  Moissan  was  the  first  to  solve  this  problem  as  far  as 
small  diamonds  are  concerned.  In  1893  he  prepared  diamonds 
in  connection  with  his  beautiful  investigations  at  very  high 
temperatures  obtained  by  means  of  the  electric  furnace.  The 
largest  diamond  which  Moissan  has  thus  far  prepared  has  a 
diameter  of  only  0.5  mm. 

The  preparation  of  large  diamonds  artificially  is  as  yet  an 
unsolved  problem. 

Another  crystalline  modification  of  carbon  is  known  as 
graphite  or  plumbago.  While  the  diamond  is  comparatively  rare, 
graphite  occurs  in  nature  in  considerable  quantities,  especially 

189 


190  ELEMENTS   OF   INOKGANIC   CHEMISTRY 

in  Siberia.  Graphite  can  readily  be  prepared  by  heating 
amorphous  carbon,  such  as  ordinary  charcoal,  in  an  electric 
furnace,  or  better,  by  dissolving  carbon  in  molten  metals  and 
allowing  it  to  crystallize. 

Amorphous  Forms  of  Carbon.  —  Carbon  occurs  in  the  uncry stal- 
lized  condition  in  many  forms.  One  of  the  best  known  is 
charcoal)  or  wood  charcoal  as  it  is  usually  termed.  If  a  piece 
of  wood  is  heated  to  a  high  temperature  in  the  presence  of  an 
abundance  of  oxygen,  the  carbon  unites  with  the  oxygen,  form- 
ing the  well-known  compound,  carbon  dioxide.  If,  on  the  other 
hand,  wood  is  heated  without  free  access  of  air,  many  products 
are  formed,  but  the  carbon  remains  behind  for  the  most  part  as 
charcoal. 

Bone-black  is  another  form  of  amorphous  carbon  obtained  by 
the  destructive  distillation  of  bones.  To  obtain  it  in  pure  form 
the  inorganic  matter  contained  in  the  bones  is  dissolved  out  by 
some  strong  acid. 

Soot  or  lamp-black  is  an  amorphous  form  of  carbon  obtained 
by  introducing  a  cold  object  into  the  flame  of  an  ordinary  lamp. 
Under  these  conditions  some  of  the  carbon,  before  it  combines 
with  oxygen,  is  deposited  in  a  very  finely  divided  condition 
known  as  lamp-black  or  soot. 

Coal  or  stone-coal  is  the  form  in  which  free  carbon  occurs 
most  abundantly  in  nature.  There  are  great  beds  of  these 
deposits  in  many  places  on  the  earth,  and  these  are  of  funda- 
mental importance  for  the  welfare  of  the  human  race.  In  coal 
we  find  vast  quantities  of  intrinsic  energy  which  can  readily  be 
converted  into  other  forms,  and  our  steam  engines,  electric 
motors,  electric  light  plants,  etc.,  are  all  dependent  upon  coal 
for  their  utility. 

These  deposits  of  coal  are  chiefly  of  vegetable  origin.  In 
certain  localities,  where  there  has  been  a  great  accumulation  of 
vegetable  matter,  this  has  undergone  decomposition  without 
free  access  of  air  and  the  carbon  has  been  deposited  in  the 
form  of  coal. 

If  the  coal  is  hard  and  comparatively  free  from  volatile  oils, 


CAEBON,    SILICON,   BORON  191 

it  is  called  anthracite;  if  it  is  soft  and  contains  much  volatile 
matter,  it  is  known  as  bituminous  coal.  If  the  process  of  coal 
formation  is  not  very  far  advanced,  we  have  peat,  lignite,  etc. 

The  Different  Forms  of  Carbon  contain  Different  Amounts 
of  Energy.  —  It  is  obvious  from  the  above  that  many  forms  of 
carbon  are  known.  The  question  arises,  How  do  these  forms 
differ  from  one  another  ?  They  are  all  carbon,  and  materially 
considered  nothing  but  carbon,  and  yet  have  very  different 
properties.  We  have  met  with  analogous  cases  in  the  different 
modifications  of  oxygen,  sulphur,  and  phosphorus,  and  found 
in  every  one  of  these  cases  that  the  different  modifications  of 
the  same  element  contain  different  amounts  of  intrinsic  energy. 
We  would  naturally  look  for  the  same  differences  in  the  case 
of  carbon. 

The  heats  of  combustion  of  different  modifications  of  carbon 
were  determined  and  the  following  results  obtained :  — 


HEAT  OF  COMBUSTION 


Charcoal 
Retort  carbon 
Diamond 
Graphite 


96.980  calories 
96.530  calories 
94.550  calories 
93.360  calories 


Since  the  end  product  is  the  same  in  every  case, — carbon 
dioxide,  —  the  differences  between  the  heats  of  combustion  of 
the  various  forms  of  carbon  are  a  measure  of  the  different 
amounts  of  intrinsic  energy  in  these  different  forms. 

We  see  that  these  differences  are  quite  considerable,  amor- 
phous carbon  having  the  largest  amount  of  intrinsic  energy 
and  the  crystallized  varieties  the  least. 

Physical  Properties  of  Carbon.  —  Carbon,  except  in  the  form 
of  the  diamond,  is  a  black  solid,  hard,  and  having  a  more  or 


192  ELEMENTS   OF  INORGANIC   CHEMISTRY 

less  metallic  lustre  in  graphite  and  anthracite,  soft  in  wood 
charcoal  and  coke,  and  a  fine  powder  in  soot  or  lamp-black. 

Carbon  remains  solid  until  an  enormously  high  temperature 
is  reached.  In  the  electric  arc,  where  the  temperature  is 
probably  in  the  neighborhood  of  3500°,  carbon  vaporizes,  but 
even  at  this  enormously  high  temperature,  comparatively 
slowly. 

COMPOUNDS   OF   CARBON 

Compounds  of  carbon  with  hydrogen,  oxygen,  nitrogen,  sul- 
phur, etc.,  are  known  in  such  large  numbers  that  a  separate 
branch  of  chemistry  has  grown  up  around  the  element  carbon. 

While  the  study  of  the  compounds  of  carbon  belongs  to 
organic  chemistry,  we  shall  take  up  a  few  typical,  fundamental 
substances,  to  give  an  idea  of  the  kind  of  compounds  which 
carbon  forms  with  other  elements. 

Compounds  of  Carbon  with  Hydrogen.  —  Carbon  forms  with 
hydrogen  a  very  large  number  of  compounds.  Indeed,  these 
two  elements  form  several  series  of  compounds ;  the  individual 
members  of  any  series  differing  in  composition  by  one  carbon 
atom  and  two  hydrogen  atoms.  The  simplest  compound  of 
carbon  and  hydrogen  has  the  composition  CH4,  and  is  known 
as  marsh  gas,  or  methane. 

There  is  a  series  of  compounds  closely  related  to  methane, 
and  known  as  the  methane  series.  The  simpler  members  are  — 

Methane CH4 

Ethane C2H6 

Propane  . C3H8 

Butane C4H10 

Carbon  forms  with  hydrogen  a  compound  containing  just 
twice  as  much  carbon  in  proportion  to  hydrogen  as  methane. 
This  compound  has  the  composition  C2H4,  and  is  known  as 
ethylene.  This,  like  methane,  is  a  fundamental  substance,  and 
the  first  member  of  a  group  of  hydrocarbons  which  have  a  con- 
stant difference  in  composition. 


CARBON,   SILICON,   BORON  193 

The  first  few  members  of  this  group  are :  — 

Ethylene C2H4 

Propylene C3H6 

Butylene C4H8 

Such  series  of  compounds  as  the  above,  in  which  successive 
members  differ  in  composition  by  the  group  CH2,  are  known  as 
homologous  series  of  compounds. 

The  above  series,  of  which  several  members  are  known,  is 
the  ethylene  series  of  hydrocarbons. 

Carbon  forms  other  series  of  compounds  with  hydrogen,  but 
it  would  lead  us  too  far  to  discuss  them  in  this  connection. 
The  acetylene  hydrocarbons,  the  first  member  of  which  is 
acetylene,  C2H2,  should,  however,  be  mentioned. 

Compounds  of  Carbon  with  Oxygen.  —  Carbon  forms  two  com- 
pounds with  oxygen,  —  carbon  monoxide,  CO,  and  carbon  diox- 
ide, C02. 

Carbon  Monoxide,  CO.  —  Carbon  monoxide  is  formed  by  the 
direct  union  of  the  two  elements.  When  carbon  is  heated  in  a 
limited  supply  of  oxygen,  the  product  is  carbon  monoxide; 

The  most  convenient  method  of  preparing  carbon  monoxide 
is  by  heating  formic  acid,  a  compound  having  the  composition 
H2C02,  with  concentrated  sulphuric  acid.  This  decomposes  as 
follows :  — 

H2C02=H20  +  CO. 

Carbon  monoxide  is  a  colorless,  poisonous  gas.  When 
breathed  into  the  system  it  combines  with  the  haemoglobin  of 
the  blood  and  prevents  the  latter  from  carrying  out  its  normal 
function,  which  is  to  carry  oxygen  to  the  various  organs  of  the 
body. 

Carbon  monoxide  is  frequently  formed  in  the  incomplete  com- 
bustion of  carbon  in  poorly  ventilated  furnaces.  From  such 
furnaces  it  easily  escapes  into  the  room,  and  is  breathed  by  the 
inhabitants.  If  the  room  into  which  carbon  monoxide  is 
escaping  is  poorly  ventilated,  bad  results  may  follow. 


194  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Carbon  monoxide  combines  directly  with  oxygen,  forming 
carbon  dioxide,  and  is,  therefore,  a  good  reducing  agent.  When 
carbon  monoxide  is  brought  in  contact  with  the  hot  oxides 
of  the  heavy  metals,  they  are  reduced  to  the  metallic  condition, 
and  the  carbon  monoxide  is  oxidized  to  carbon  dioxide. 

Thermochemistry  of  Carbon  Monoxide.  —  When  a  grain- 
atomic  weight  (12  grams)  of  carbon  is  burned  to  carbon  mo- 
noxide, the  amount  of  heat  set  free  is  only  about  2000  calories, 
while  6000  calories  are  liberated  when  carbon  monoxide  is 
oxidized  to  carbon  dioxide.  Carbon  monoxide,  therefore,  con- 
tains a  large  amount  of  intrinsic  energy  which  can  be  converted 
into  heat  by  simply  oxidizing  it  to  carbon  dioxide.  It  is  due 
to  this  fact  that  carbon  monoxide  is  an  excellent  heating  agent. 

Water-gas.  —  One  of  the  methods  for  preparing  carbon 
monoxide  is  to  pass  water-vapor  over  highly  heated  carbon; 
the  action  which  takes  place  being  — 

C  +  H20  =  CO  +  H2. 

This  mixture  of  carbon  monoxide  and  hydrogen  would  have 
very  little  value  as  an  illuminating  gas,  since  both  of  these 
gases  burn  with  a  comparatively  colorless  flame,  although  they 
evolve  an  enormous  amount  of  heat.  This  mixture  of  gases  is 
passed  through  highly  heated  petroleum  vapor,  and  is  thus 
mixed  with  hydrocarbons  and  other  substances  which  give  oft' 
an  abundance  of  light  when  they  are  burned.  This  mixture, 
known  as  water-gas,  is  now  used  largely  for  illuminating 
purposes. 

Water-gas  is  now  extensively  used  where  illuminating  gas, 
made  by  the  dry  distillation  of  coal,  was  formerly  employed. 

Carbon  Dioxide,  C02.  —  The  highest  product  of  the  direct 
oxidation  of  carbon  is  carbon  dioxide.  This  compound  occurs 
in  a  number  of  places  in  the  free  condition.  It  is  one  of  the 
constituents,  as  will  be  remembered,  of  atmospheric  air.  It 
also  occurs  dissolved,  in  greater  or  less  quantity,  in  water. 

Carbon  dioxide  escapes  from  certain  mineral  springs  dissolved 
in  the  water  of  such  springs. 


CARBON,    SILICON,   BORON  195 

When  the  water  reaches  the  surface  of  the  earth,  a  part  of 
the  gas  escapes  and  gives  the  characteristic  effervescence. 

Carbon  dioxide  is  given  off  by  animals,  as  can  be  readily 
shown  by  breathing  for  a  short  time  into  lime-water  or  a  solu- 
tion of  oarium  hydroxide,  when  insoluble  calcium  or  barium 
carbonate  is  formed. 

Carbon  dioxide  is  also  set  free  when  animal  and  vegetable 
matter  decomposes,  and  also  in  many  processes  of  fermenta- 
tion. 

Preparation  of  Carbon  Dioxide.  —  Carbon  dioxide  can  be  pre- 
pared by  a  number  of  methods.  Theoretically,  one  of  the 
simplest  methods  is  to  burn  carbon  in  an  excess  of  air :  — 

C  +  02=C02. 

Practically,  a  far  more  convenient  method  of  preparing 
carbon  dioxide  is  to  treat  a  carbonate  with  an  acid:  — 

K2C03  +  2  HC1  =  2  KC1  +  H2C03. 

The  hypothetical  compound  H2C03  breaks  down  thus :  — 
H2C03  =  H20  +  C02. 

When  certain  compounds  are  heated,  they  readily  lose  carbon 
dioxide,  while  others  lose  it  only  at  high  temperatures.  The 
carbonate  of  calcium,  or  ordinary  limestone  or  marble,  belongs 
to  the  former  class.  When  this  substance  is  heated,  it  breaks 
down  thus :  — 

CaC03  =  CaO  +  C02. 

Chemical  Properties  of  Carbon  Dioxide.  —  The  most  character- 
istic chemical  property  of  carbon  dioxide  is  its  power  to  form 
salts  in  the  presence  of  aqueous  solutions  of  alkalies.  When 
one  equivalent  of  caustic  potash,  in  water,  is  brought  into  the 
presence  of  one  equivalent  of  carbon  dioxide,  the  following 
reaction  takes  place  :  — 

KOH  +  C02  =  KHC03. 


196  ELEMENTS   OF   INORGANIC   CHEMISTRY 

When  two  equivalents  of  caustic  potash  are  used,  we  have 
the  following  reaction  :  — 

2  KOH  +  C02  =  K2C03  +  H20. 

The  first  salt  is  acid  potassium  carbonate,  the  second  normal 
potassium  carbonate. 

Carbon  dioxide  in  the  presence  of  water  acts,  then,  as  a 
dibasic  acid. 

The  aqueous  solution  reacts  slightly  acid,  showing  that  there 
are  a  small  number  of  hydrogen  ions  present  :  — 

H2C03  =  H,  HC03. 

The  acid  is  so  weak  that  we  are  not  justified  in  assuming 
that  there  is  any  dissociation  of  the  ion,  HC03,  in  the  pres- 
ence of  water  alone.  When  an  alkali  is  present,  and  all  the 
hydrogen  ions  from  the  first  stage  of  dissociation  are  used  up, 

it  is  probable  that  the  ion,  HC03,  begins  to  dissociate  thus  :  — 

HC03  =  H  +  C03, 

and  this  dissociation  continues  to  the  end  if  there  are  enough 
hydroxyl  ions  from  the  base  present  to  combine  with  all  the 
hydrogen  ions  as  rapidly  as  they  are  formed. 

The  carbonates,  like  the  salts  of  all  weak  acids,  are  hydro- 
lyzed  (see  page  176)  by  water.  This  is  shown  by  the  fact  that  a 
salt  like  sodium  carbonate  shows  a  strongly  alkaline  reaction, 
which  means  that  there  are  hydroxyl  ions  present  :  — 


+  H20  =  Na,  OH  +  Na,  HC03. 

The  hydrolysis  of  carbonates  is  by  no  means  complete,  only 
a  comparatively  small  number  of  molecules  being  broken  down 
by  the  water  as  represented  in  the  above  equation. 

REDUCTION  OF   CARBON  DIOXIDE   BY  PLANTS 

Carbon  dioxide  is  being  continually  reduced  by  green  plants 
in  the  sunlight.  They  build  up  the  carbon  into  complex  com- 


CAKBON,    SILICON,   BOKON  197 

pounds  with  hydrogen,  oxygen,  and  perhaps  nitrogen,  and 
these  compounds  contain  enormous  amounts  of  intrinsic  energy. 
The  complex  compounds  of  carbon  are  consumed  by  animals, 
which  decompose  these  substances  into  much  simpler  ones, 
especially  into  carbon  dioxide.  The  large  excess  of  intrinsic 
energy  in  the  complex  compounds  over  that  in  the  simpler 
products  which  animals  excrete,  is  converted  into  heat  and  by 
the  animal  into  mechanical  work. 

The  chief  source  of  the  energy  which  animals  expend  is, 
then,  the  complex  compounds  of  carbon,  which  are  built  up  by 
plants  from  carbon  dioxide. 

It  is  of  interest  to  note  that  most  of  the  carbon  in  animal 
and  vegetable  tissues  ultimately  passes  off  when  these  decay, 
in  the  form  of  carbon  dioxide.  The  carbon  dioxide  is  again 
taken  up  by  the  plant,  converted  into  complex  compounds, 
consumed  by  the  animal,  broken  down  into  simpler  substances, 
and  the  cycle  is  thus  completed. 

Physical  Properties  of  Carbon  Dioxide.  —  The  gas  carbon 
dioxide  can  be  readily  liquefied.  Carbon  dioxide  is  liquefied 
on  a  large  scale  by  pumping  it  into  thick-walled,  steel  cylin- 
ders, which  are  kept  cool.  Such  cylinders  are  kept  in  the 
laboratory  as  sources  of  supply  of  carbon  dioxide. 

When  the  carbon  dioxide  is  allowed  to  escape  from  such 
cylinders  through  a  fine  opening,  part  of  it  volatilizes  and 
escapes  as  a  gas,  while  the  remainder  is  converted  into  the 
solid  condition  and  can  be  caught  in  a  bag  placed  over  the 
jet.  Solid  carbon  dioxide  is  a  compact  white  mass  resembling 
compressed  snow.  It  has  been  extensively  used  as  a  refrigerat- 
ing agent. 

Compounds  of  Carbon  with  Oxygen  and  Hydrogen.  —  Thou- 
sands of  such  compounds  are  known.  While  these  belong 
strictly  to  the  subject  of  organic  chemistry,  a  few  typical  sub- 
stances will  be  considered  here. 

The  alcohols  are  among  the  simplest  of  the  compounds  of 
carbon  with  oxygen  and  hydrogen.  The  first  member  of  this 
series  of  compounds  is  methyl  alcohol,  CH40,  or  wood  spirit, 


198  ELEMENTS   OF   INORGANIC   CHEMISTRY 

as  it  is  termed.  The  alcohols  form  a  homologous  series  of 
compounds  which  are  analogous  to  the  hydrocarbons.  The 
first  members  of  this  series  are :  — 

Methyl  alcohol CH40 

Ethyl  alcohol C2H60 

Propyl  alcohol C3H8O 

A  product  of  the  oxidation  of  an  alcohol  is  an  ether.  Take 
ordinary  ethyl  ether.  This  has  the  composition  C4H100,  and 
is  a  member  of  a  homologous  series  of  compounds.  The  first 
member  is  methyl  ether,  C2H60,  the  second  methyl-ethyl  ether 
C3H80,  and  so  on. 

If  an  alcohol  is  further  oxidized,  it  passes  over  into  an  acid, 
and  we  have  homologous  series  of  organic  acids.  If  ethyl 
alcohol  is  properly  oxidized,  we  have  acetic  acid,  CH3COOH. 
This  is  the  second  member  of  the  formic  acid  series,  formic 
acid  being  the  first. 

Formic  acid HCOOH 

Acetic  acid  CH3COOH 

Propionic  acid C2H5COOH 

Compound  of  Carbon  with  Sulphur,  CS2.  —  Carbon  bisulphide, 
CS2,  is  formed  by  passing  the  vapors  of  sulphur  over  highly 
heated  carbon.  The  two  elements  unite,  forming  carbon  bisul- 
phide, which,  being  volatile,  passes  out  of  the  field  of  action. 
Carbon  bisulphide  is  very  easily  inflammable,  readily  uniting 
with  oxygen  and  forming  carbon  dioxide  and  sulphur  dioxide. 

Carbon  bisulphide  is  an  excellent  solvent,  not  only  for  oils, 
fats,  and  other  complex  organic  compounds,  but  for  bromine 
and  iodine  and  many  other  substances.  It  is  a  liquid  with  a 
highly  disagreeable  odor,  boiling  at  46°. 

Compound  of  Carbon  with  Nitrogen  —  Cyanogen  (CN)2. 
— When  we  consider  the  inert  nature  of  the  element  nitrogen, 
it  is  surprising  that  the  compound  (CN)2  should  be  capable  of 
existence.  The  two  elements,  however,  do  not  combine 
directly,  but  combine  readily  with  an  alkali  metal,  forming 


CARBON,    SILICON,   BORON  199 

such  compounds  as  potassium  cyanide,  KCN.  Cyanogen  is 
not  so  readily  obtained  from  the  potassium  compound,  but  is 
very  easily  prepared  from  mercuric  cyanide.  This  compound 
when  heated  breaks  down  into  mercury  and  cyanogen  :  — 


Cyanogen  is  a  gas  which  is  characterized  by  its  extremely 
poisonous  nature. 

Hydrocyanic  Acid,  HCN.  —  Hydrocyanic  acid  is  formed  when 
a  cyanide  is  treated  with  an  acid  :  — 

MCN  +  HC1  =  MCI  +  HCK 

This  acid,  which  is  very  soluble  in  water,  is  known  as  prussic 
acid.  It  is  characterized  by  its  extremely  poisonous  nature. 
Hydrocyanic  acid  is  a  very  weak  acid.  It  dissociates  into 

4-          - 

H,  ON,  but  only  to  a  slight  extent.  When  there  is  a  strong 
alkali  present,  which  is  the  same  as  to  say  an  excess  of 
hydroxyl  ions,  the  hydrogen  ions  are  used  up  as  fast  as  they 
are  formed,  combining  with  the  hydroxyl  ions  to  form  water, 
and  more  of  the  acid  dissociates. 

THE  ROLE  OF  CARBON  IN  PRODUCING  LIGHT 

Illumination.  —  The  subject  of  the  production  of  light  or 
illumination  is  one  that  has  attracted  attention  for  a  very  long 
time.  In  practically  all  of  the  earlier  methods  of  producing 
light,  and  in  many  of  those  used  to-day,  carbon  is  employed 
in  one  form  or  another.  Most  of  the  methods  of  illumination 
owe  their  existence  to  some  compound  of  carbon  which  is 
burned  or  oxidized,  giving  out  heat  and  light. 

Candle  and  Oil-lamp.  —  In  the  candle  we  have  complex  com- 
pounds of  carbon  which  at  ordinary  temperatures  are  solid. 
These  are  made  by  melting  the  stearine,  tallow,  or  paraffine, 

id  pouring  the  liquid  into  a  mould  after  a  wick  has  been 
placed  in  the  centre  of  the  mould.  The  object  of  the  wick  is 


200 


ELEMENTS   OF  INORGANIC   CHEMISTRY 


-—b 


to  carry  by  capillarity  the  material  of  the  candle,  as  it  is 
melted,  up  into  the  flame.  The  tip  of  the  candle  is  melted, 
and  the  end  of  the  wick  ignited.  This  melts  a  portion  of  the 
solid  hydrocarbons,  which  are  carried  up  by  the  wick  into  the 
flame  and  are  burned,  the  heat  set  free  melting  more  of 
the  solid,  and  the  process  is  thus  a  continuous  one.  The  heat 
and  light  are  derived  from  the  breaking  down  of  complex 
compounds  of  carbon  into  simpler  substances,  and  the  oxidation 
of  the  carbon  to  carbon  dioxide. 

In  our  oil-lamps  the  compounds  of  carbon  which  are  to  be 
burned  are  liquid  at  ordinary  temperatures.     These  are  carried 
up  into  the  flame  by  means  of  the  wick,  as  in 
the  candle,  and  the  same  general  processes  are 
involved  in  the  production  of  light  and  heat. 

Flames  and  their  Luminosity. — If  we  examine 
a  typical  flame,  say  that  of  a  candle,  we  observe 
three  distinct  parts  :  An  inner  cone  a  (Fig.  50), 
of  unburned  gases,  is  surrounded  by  a  zone  b, 
of  partially  oxidized  substances.  It  is  in  this 
zone  that  acetylene,  which  has  much  to  do  with 
the  light-giving  power  of  the  flame,  is  formed. 
This  zone  is  the  chief  source  of  light  in  the 
flame.  This  is  surrounded  by  a  third  layer  c, 
of  burning  gases,  and  here,  where  there  is  an 
abundant  supply  of  oxygen  from  the  air,  the 
processes  of  oxidation  are  completed.  This 
part  of  the  flame  is  relatively  only  slightly 
luminous." 

So  much  for  the  structure  of  a  flame.  The 
question  remains,  What  are  the  causes  of  the 
luminosity  of  flames  ?  We  have  seen  that  the  chief  source  of 
light  in  a  flame  is  in  the  middle  zone,  where  the  combustion  is 
not  complete.  This  gave  rise  to  the  theory  that  the  chief  source 
of  light  in  a  flame  is  unburned,  solid  particles  of  carbon,  which 
become  heated  to  incandescence.  These  particles  come  from 
the  compounds  of  carbon  which  are  decomposed  by  the  heat  of 


--a 


FIG.  50. 


CARBON,    SILICON,   BORON 


201 


the  flame.     This  theory  accounts  for  many  of  the  facts  concern- 
ing the  luminosity  of  flames,  but  by  no  means  for  all. 

Hydrogen  gas  at  atmospheric  pressure  burns  with  an  almost 
non-luminous  flame.  Hydrogen  gas,  under  high  pressure,  how- 
ever, burns  with  a  luminous  flame.  The  effect  of  pressure  on 
luminosity  is  also  shown  by  burning  a  candle  in  a  valley  and 
on  a  high  mountain.  Under  the  former  conditions,  where  the  ' 
pressure  is  relatively  high,  the  luminosity  is  much  greater. 
Further,  gases  which  burn  with  a  luminous  flame  can  be  made 
to  burn  with  a  relatively  non-luminous  flame  by  mixing  them 
with  an  indifferent  gas,  or  by  simply  cooling  them. 

These  facts  cannot  by  any  means  be  all  explained  on  the  solid 
particle  theory  of  luminosity.  There  are  undoubtedly  many 
influences  which  affect  the  luminosity  of  flames,  so  that  prob- 
ably no  one  theory  can  account 
for  all  of  the  facts  connected 
with  this  phenomenon. 

According  to  recent  investi- 
gations it  seems  very  probable 
that  the  formation  and  oxida- 
tion of  acetylene  in  flames  is 
vitally  connected  with  their 
light-giving  power. 

Bunsen  Burner ;  Blowpipe. 
—  Practical  use  is  made  of  the 
fact  that  complete  oxidation, 
and  also  the  dilution  of  a  gas 
with  an  indifferent  gas,  lowers 
its  luminosity,  in  constructing 
the  bunsen  burner. 

A  buusen  burner  is  shown 
in  Figure  51.  The  gas  enters 
through  the  horizontal  tube  A,  into  the  vertical  tube  B.  Air 
enters  through  the  hole  (7,  and  mixes  with  the  gas.  The  flame 
consists  of  two  distinct  parts :  an  inner  blue  cone,  where  the 
oxidation  is  far  from  complete,  and  which  is  known  as  the, 


FIG.  51. 


202 


ELEMENTS  OF   INORGANIC   CHEMISTRY 


reducing  flame,  since  it  lias  remarkable  power  to  combine  with 
oxygen  and  reduce  substances  such  as  the  oxides  of  the  metals  ; 
and  an  outer,  almost  non-luminous  tip,  where  the  oxidation  of 
the  gas  is  completed  and  where  the  temperature  is  very  high, 
This  is  known  as  the  oxidizing  flame,  on  account  of  its  power 
to  give  up  oxygen  to  substances  which  can  be  oxidized.  Metals, 
for  example,  in  this  flame  are  usually  converted  into  oxides. 

By  means  of  the  bunsen  burner  very  high  temperatures  can 
be  secured  by  the  combustion  of  illuminating  gas, 
without  the  production  of  any  appreciable  quantity 
of  light. 

When  a  cold  object  is  inserted  into  the  flame 
of  a  bunsen  burner,  no  carbon  is  deposited  upon  it 
in  the  form  of  soot,  and  the  lamp  is,  therefore,  very 
convenient  for  heating  where 
cleanliness  is  absolutely  essen- 
tial.    The  bunsen  burner  is  one 
of  the   most   frequently  used 
pieces  of  appara- 
tus in  the  chemical 
laboratory. 

The  blowpipe 
is  a  still  more 
efficient  means 
of  obtaining  a 
clean  oxidizing 
and  a  clean  reduc- 
ing flame,  and 
of  directing  these 
flames  where  they 
are  desired.  The  blowpipe  itself  is  shown  in  Figure  52.  It 
consists  of  a  tube  t,  into  which  the  breath  is  blown  from  the 
mouth,  and  a  tube  f1?  at  right  angles  to  this,  through  which  the 
ajr  from  the  lungs  passes  into  the  flame.  Into  the  top  of  an 
ordinary  buusen  burner  is  inserted  a  tube  with  a  narrow  open- 
ing, so  as  to  give  a  narrow  flame. 


FIG.  52. 


CARBON,    SILICON,   BORON  203 

The  blowpipe  is  placed  upon  the  upper  edge  of  this  tube,  as 
indicated  in  the  drawing,  and  the  breath  expelled  continu- 
ously through  the  tube.  The  combustion  of  the  gases  is 
excellent,  and  the  flame  takes  the  form  shown  in  the  figure. 
The  inner  flame  a  is  the  reducing  flame,  and  the  outer  tip  b 
the  oxidizing  portion  of  the  flame. 

By  means  of  the  blowpipe  flame  very  delicate  work  can  be 
done.  In  the  reducing  flame  small  quantities  of  metal  oxides 
can  be  reduced  to  the  metallic  condition,  and  identified. 

Effect  of  Cooling  the  Flame.  —  The  effect  of  cooling  the  flame 
can  be  readily  shown  by  means  of  the  following  experiment : 
Open  an  ordinary  gas  stop-cock,  and  a  short  distance  above  the 
orifice  hold  a  wire  gauze  with  fine  mesh.  Light  the  gas  above 
the  gauze,  and  it  will  burn  without  the  gas  below  the  gauze 
taking  fire.  This  is  due  to  the  fact  that  the  metallic  gauze  con- 
ducts the  heat  away  so  rapidly,  that  the  gas  below  the  gauze  is 
not  heated  to  its  kindling  temperature  and  does  not  ignite. 

This  principle  was  made  use  of  by  Sir  Humphry  Davy  in  the 
construction  of  his  safety  lamp  for  use  in  mines  where  explosive 
gases  are  liable  to  accumulate.  The  flame  is  simply  surrounded 
by  a  fine  wire  gauze.  If  the  explosive  gases  should  ignite  on 
the  inside  of  the  gauze,  the  flame  cannot  propagate  itself 
through  the  gauze,  since  it  is  too  greatly  cooled.  The  gauze, 
conducting  the  heat  from  the  flame,  prevents  the  gases  on 
the  outside  from  becoming  heated  to  their  kindling  tempera- 
ture, and  thus  explosions  are  avoided  when  lights  are  carried 
into  an  atmosphere  containing  explosive  gases. 

The  Welsbach  Light.  —  The  Welsbach  light  differs  from  the 
ordinary  gas-light  in  that  solid  substances  are  introduced  into 
the  flame,  which,  when  hot,  have  remarkable  light-giving  power. 
The  Welsbach  light  depends  for  its  value  entirely  upon  the 
mantle.  The  mantle  consists  of  a  mixture  of  thorium  and 
cerium  oxides.  It  is  prepared  as  follows :  Fine  cotton  thread 
is  woven  into  exactly  the  form  of  the  mantle.  This  is  satu- 
rated with  an  aqueous  solution  of  a  mixture  of  the  nitrates  of 
thorium  and  cerium.  This  mixture  contains  99  per  cent  of  the 


204  ELEMENTS   OF  INOKGANIC    CHEMISTRY 

thorium  salt  and  one  per  cent  of  the  cerium  salt.  The  mantle 
is  then  dried  and  highly  heated  to  burn  out  all  organic  matter, 
and  to  convert  the  cerium  and  thorium  nitrates  into  the  oxides. 
It  is  then  ready  for  use. 

It  is  a  remarkable  fact  that  if  the  amount  of  cerium  salt 
added  to  the  thorium  is  either  increased  or  diminished  appreci- 
ably, the  light-giving  power  of  the  Welsbach  burner  is  greatly 
diminished. 

The  value  of  the  burner  is  to  be  found  in  the  power  of  these 
oxides  to  convert  heat  energy  in  large  quantity  into  light 
energy,  so  that  the  final  result  is  a  conversion  of  more  of  the 
intrinsic  energy  of  the  carbon  compounds  and  other  substances 
in  the  gas  into  light  energy. 

The  Electric  Light. —At  first  sight  the  relation  between 
carbon  and  the  electric  light  may  not  appear  to  be  very  close, 
other  than  the  use  of  carbon  as  the  source  of  energy  to  drive 
the  dynamo  which  generates  the  electrical  energy. 

To  obtain  light  energy  from  electrical  energy,  a  resistance  to 
the  passage  of  the  electrical  current  is  interposed.  The  current 
is  usually  passed  between  two  carbon  poles,  which  are  heated 
to  such  a  high  temperature  that  the  carbon  is  partially  volatil- 
ized. At  this  temperature  the  highly  heated  carbon  gives  off 
an  enormous  amount  of  light  energy,  and  this  is  the  source  of 
the  light  in  the  electric  arc-light.  In  the  incandescent  light 
the  carbon  is  heated  white-hot  in  a  vacuum  or  an  inert  gas, 
and  gives  out  light  without  undergoing  any  appreciable  chemical 
change. 

SILICON  (At.  Wt.  =  28.4) 

The  second  member  of  the  fourth  group  in  the  Periodic 
System  is  silicon.  This  element  is  very  widely  distributed 
over  the  surface  of  the  earth,  and  constitutes  an  important  part 
of  most  rocks.  Silicon  occurs  in  great  abundance  as  the  dioxide, 
and  forms  an  acid  —  silicic  acid,  whose  salts  make  up  many  of 
our  best-known  rocks.  Silicon  dioxide,  or  quartz,  also  occurs 
in  huge  masses,  and  is  a  constituent  of  many  rocks?  especially 


CARBON,    SILICON,   BORON  205 

granites,  gneisses,  etc.    Silicon  dioxide  occurs  in  great  abundance 
as  sand,  especially  along  the  edges  of  large  bodies  of  water. 

Preparation  of  Silicon.  —  Silicon  is  prepared  from  its  com- 
pounds by  a  number  of  methods.  One  of  these  consists  in 
heating  silicon  tetrafluoride  with  sodium:  — 


Silicon  Hydride  or  Hydrogen  Silicide,  SiH4.  —  Silicon  forms 
a  compound  with  hydrogen,  known  as  silicon  hydride,  or  hy- 
drogen silicide,  containing  one  atom  of  silicon  united  with  four 
atoms  of  hydrogen.  It  also  forms  the  compound  Si2H6,  which 
is  spontaneously  inflammable. 

As  far  as  composition  is  concerned,  silicon  tetrahydride  is 
analogous  to  methane,  —  SiH4.  —  CH4.  The  former,  however, 
is  very  unstable,  while  the  latter  is  quite  stable. 

Silicon  Dioxide,  Si02.  —  Silicon  forms  one  compound  with 
oxygen,  —  silicon  dioxide.  This  is  analogous  to  carbon  dioxide. 
It  does  not  form  the  analogue  of  carbon  monoxide. 

Silicon  dioxide  occurs  in  nature  in  great  abundance.  It  is 
beautifully  crystalline  in  several  varieties  of  quartz,  such  as 
amethyst,  rock  crystal,  and  the  like,  and,  with  certain  impurities 
which  give  it  color,  it  is  of  more  or  less  value  as  gems,  such  as 
opal,  jasper,  onyx,  agate,  etc. 

It  occurs  in  great  masses  in  less  attractive  forms,  such  as 
quartz,  sand,  flint,  sandstone,  and  the  like,  and  is  frequently  the 
chief  constituent  of  large  mountain  ranges.  When  we  con- 
sider the  abundance  of  the  two  forms,  quartz  and  sandstone, 
we  can  see  the  importance  of  the  element  silicon  in  the  inor- 
ganic world,  and  from  a  geological  standpoint. 

Silicon  dioxide  is  very  resistant  to  chemical  reagents,  and  is 
not  attacked  by  acids,  with  the  exception  of  hydrofluoric. 
When  powdered  very  finely  and  fused  with  a  caustic  alkali,  or 
an  alkaline  carbonate,  it  is  transformed  into  a  silicate:  — 

Si02  +  2  KOH  =  K2Si03  +  H20. 
This  is  a  salt  of  a  silicic  acid  having  the  composition  H2Si03. 


206  ELEMENTS  OF   INORGANIC   CHEMISTRY 

The  Acids  of  Silicon.  —  Silicon  combines  with  hydrogen  and 
oxygen,  forming  a  number  of  acids  which,  however,  can  all  be 
regarded  as  derived  from  one  mother-substance,  —  normal  silicic 
acid.  When  an  alkaline  silicate  like  that  mentioned  above  is 
treated  with  an  acid,  the  following  reaction  probably  takes 
place :  — 

K2Si03  +  2  HC1  +  H20  =  2  KC1  +  H4SiO4. 

The  compound  H4Si04  is  known  as  normal  silicic  acid,  or 
orthosilicic  acid. 

When  the  normal  acid  is  heated,  it  loses  water  and  forms 
silicon  dioxide :  — 

H4Si04  =  2  H20  +  Si02. 

Polysilicic  acids  can  all  be  regarded  as  derived  from  the  acid 
H4Si04,  by  removal  of  one  or  more  molecules  of  water  from  two 
or  more  molecules  of  the  acid.  Thus,  by  the  removal  of  one 
molecule  of  water  from  two  molecules  of  normal  silicic  acid, 
we  have : — 

2  H4Si04  =  H20  +  H6Si207. 

By  removing  two  molecules  of  water :  — 

2  H4Si04  =  2  H20  +  H4Si206. 

This  series  of  acids,  salts  of  some  of  which  are  known,  sug- 
gest the  homologous  compounds  of  carbon.  The  constant  dif- 
ference with  carbon  is  the  group  CH2.  The  constant  difference 
with  silicon  is  the  molecule  of  water  H20. 

Some  of  the  salts  of  the  polysilicic  acids  are  very  important 
substances,  since  they  constitute  many  of  the  most  abundant 
silicates. 

Conversion  of  Silicates  into  Carbonates. — Notwithstanding 
the  great  stability  and  insolubility  of  the  silicates,  they  are 
being  decomposed  all  over  the  surface  of  the  earth  by  such  a 
weak  acid  as  carbonic  acid.  This  carbonation  is  taking  place 
all  over  the  surface  of  the  earth,  wherever  the  carbon  dioxide 
in  the  air  and  in  the  waters  comes  in  contact  with  silicates. 
This,  at  first  sight,  is  very  surprising.  How  is  it  possible  for 


CARBON,    SILICON,   BORON  207 

such  a  weak  acid  as  carbonic  acid  to  displace  silicic  acid  from 
the  very  stable  silicates  ?  This  is  especially  difficult  to  under- 
stand when  we  consider  that  carbonic  acid  is  so  easily  volatile, 
and  therefore  escapes  from  the  field  of  chemical  action. 

The  explanation  is  to  be  found  in  the  effect  of  mass  on 
chemical  activity.  This  is  one  of  the  very  best  examples  of 
mass  action,  as  conditioning  the  direction  as  well  as  the  magni- 
tude of  chemical  action. 

The  great  amount  of  carbon  dioxide  in  the  air  and  in  the 
water,  acting  slowly  but  continually  for  a  long  period  of  time, 
effects  a  reaction  which,  in  the  laboratory,  would  be  impossible. 

The  above  process  is  of  great  geological  and  economical  im- 
portance. By  this  means  in  part,  many  of  the  most  resistant 
rocks  are  decomposed  and  the  surface  of  the  earth  greatly 
changed  in  appearance.  This  process  is  of  importance  in  that 
the  constituents  of  rocks  are  made  available  for  plants. 

Compound  of  Silicon  with  Fluorine — Silicon  Tetrafluoride, 
SiF4.  —  The  compound  of  silicon  and  fluorine,  SiF4,  is  of  special 
interest,  since  it  is  the  compound  formed  when  hydrofluoric 
acid  acts  on  glass.  Silicon  tetrafluoride  is  formed  by  the 
action  of  hydrofluoric  acid  on  silicon  dioxide.  Since  hydro- 
fluoric acid  is  prepared  most  conveniently  by  the  action  of  sul- 
phuric acid  on  calcium  fluoride,  silicon  tetrafluoride  is  prepared 
by  mixing  sand,  calcium  fluoride,  and  sulphuric  acid. 

Carborundum,  SiC.  —  A  compound  of  silicon  with  carbon 
should  be  mentioned  on  account  of  its  remarkable  composition 
and  properties.  This  is  the  compound  SiC,  known  as  carborun- 
dum. It  is  formed  by  heating  in  an  electric  furnace  a  mixture 
of  carbon  and  finely  powdered  sand.  Carborundum  is  remarka- 
bly hard,  and  is  technically  useful  on  account  of  this  property. 
It  is  very  resistant  to  chemical  reagents. 

In  this  same  group  the  rare  elements,  germaniumj  titanium, 
zirconium,  cerium,  and  thorium  occur. 

BORON  (At.  Wt.  =  11.0) 
We  pass  now  to  group  III  of  the  Periodic  System,  the  first 


208  ELEMENTS  OF  INORGANIC  CHEMISTRY 

member  of  which  is  boron.  This  is  the  only  member  of  this 
group  which  has  distinctly  acid-forming  properties.  We  shall, 
therefore,  take  up  boron  in  the  present  connection,  and  the  re- 
maining members  of  the  group  considerably  later,  when  we 
come  to  study  the  base-forming,  or  metallic  elements. 

Occurrence,  Preparation,  and  Properties.  —  The  borates,  or 
salts  of  boric  acid,  are  the  chief  source  of  the  element  boron. 
Boron  is  prepared  by  the  reduction  of  the  trioxide  of  boron. 
In  preparing  boron  the  nitrogen  of  the  air  must  be  excluded 
by  a  layer  of  borax,  since  nitrogen  combines  with  boron  at 
high  temperatures.  Boron  forms  beautiful  crystals,  which 
are  characterized  by  their  great  hardness.  They  seem  to  have 
nearly  the  same  hardness  as  the  diamond.  Boron  trioxide  is 
formed  by  burning  boron  in  oxygen  :  — 


Boron  trioxide  is  the  anhydride  of  boric  acid,  which  we  shall 
now  study. 

Boric  Acid,  H3B03.  —  Boric  acid  occurs  in  nature  in  the  free 
condition.  It  is  volatile  with  water-vapor,  arid  in  the  region 
of  certain  hot  springs,  as  in  Tuscany,  it  is  brought  to  the  sur- 
face of  the  earth  by  the  escaping  vapors. 

Boric  acid  is  soluble  in  water,  forming  beautiful  white  crys- 
tals when  the  aqueous  solution  is  evaporated.  It  is  easily  rec- 
ognized by  the  fact  that  its  alcoholic  solution  burns  with  a 
characteristic  green  flame.  If  boric  acid,  or  a  borate  treated 
with  sulphuric  acid,  is  treated  with  a  little  alcohol  and  the 
alcohol  ignited,  the  flame  appears  green  throughout  if  there  is 
much  boric  acid  present.  If  only  a  small  amount  of  boric  acid 
is  present,  the  flame  is  green  only  on  the  edges. 

The  salts  of  the  normal  boric  acid,  H3B03,  do  not  exist. 

Boric  acid  can  lose  water  and  form  an  acid  whose  salts  are 
well  known  :  — 

4  H3B03  =  5  H20  +  H2B407. 

The  acid  H2B407  is  known  as  tetraboric  add,  and  its  sodium 
salt,  Na2B407,  10H20,  is  ordinary  borax.  Borax  melts  easily, 


CARBON,    SILICON,   BORON  209 

forming  a  colorless  liquid.  This  liquid  has  the  power  of  dis- 
solving certain  metal  oxides  and  forming  with  them,  when  cold, 
glass-like  masses  which  have  characteristic  colors.  The  borax 
bead  is,  consequently,  of  importance  in  blowpipe  analysis  for 
the  detection  of  metals. 

Boron  Nitride,  BN,  is  formed  by  the  direct  union  of  nitrogen 
with  boron  at  an  elevated  temperature. 


Summary.  —  We  have  studied  thus  far  oxygen,  hydrogen, 
and  the  halogens  or  members  of  group  VII  in  the  Periodic 
System.  The  analogues  of  oxygen  —  sulphur,  selenium,  and 
tellurium  —  in  group  VI,  were  then  taken  up.  The  remaining 
members  of  group  VI  are  so  distinctly  metallic  that  they  will 
be  referred  to  with  the  metals.  The  important  members  of 
group  V  were  then  studied,  and  all  of  the  important  elements 
in  group  IV,  with  the  exception  of  the  metals,  tin  and  lead. 

We  have  begun  the  study  of  group  III  with  the  first  mem- 
ber boron,  which  is  distinctively  an  acid-forming  element. 
The  remaining  members  of  this  group,  however,  are  so  distinc- 
tively base-forming  or  metallic  that  they  will  be  studied  with 
the  metals. 

Having  completed  our  study  of  the  non-metals,  or  metalloids 
as  they  are  termed,  we  shall  now  turn  to  the  metals,  and  with 
these  we  shall  begin  with  group  I  —  the  alkalies. 


EXPERIMENTS  WITH  CARBON,  SILICON,  AND  BORON 

Experiment  115.  Comparison  of  the  Different  Forms  of  Car- 
bonr — (Graphite;  charcoal;  bone-black  or  lamp-black;  coal.) 

Compare  the  several  varieties  of  carbon ;  it  will  be  seen  that 
graphite  is  hard  and  has  crystalline  structure.  Charcoal  and 
bone-black  are  soft  and  without  any  crystalline  form.  Certain 
varieties  of  coal  are  hard,  while  others  are  quite  soft.  In 
general  the  bituminous  coals  are  soft  and  the  anthracites  hard. 
(See  page  190.) 


210 


ELEMENTS   OF   INORGANIC   CHEMISTRY 


Experiment  116.  Absorption  of  Gases  and  Coloring-matter 
by  Charcoal.  —  (Dilute  solution  of  hydrogen  sulphide ;  solution 
of  litmus;  bone-black.) 

That  charcoal  will  absorb  a  gas  like  ammonia  was  strikingly 
shown  in  Experiment  49. 

Shake  a  dilute  solution  of  hydrogen  sulphide  in  water  with 
an  excess  of  bone-black,  and  filter  the  solution.  Practically 
all  of  the  hydrogen  sulphide  will  be  absorbed  by  the  bone- 
black,  and  the  solution  will  be  nearly  free  from  the  odor  of 
the  gas. 

Shake  a  solution  of  litmus  with  a  considerable  quantity  of 
bone-black.  Filter.  If  the  solution  is  not  colorless,  boil  it  with 
more  bone-black  and  again  filter.  All  of  the  coloring-matter 
will  now  be  removed  from  the  solution  by  the  bone-black. 
(See  page  191.) 


FIG.  53. 


Experiment  117.  Reducing  Power  of  Carbon. — (Lead  oxide ; 
copper  oxide.) 

Heat  a  mixture  of  lead  oxide  and  finely  powdered  carbon  in 
a  hard-glass  tube,  and  conduct  any  gases  which  escape  into  a 


CARBON,    SILICON,   BORON  211 

solution  of  barium  hydroxide  in  another  test-tube.  Insoluble 
barium  carbonate  will  be  formed,  showing  that  the  carbon  has 
been  oxidized  to  carbon  dioxide.  That  the  zinc  oxide  has  lost 
oxygen  and  been  reduced  to  metallic  zinc,  can  be  shown  by 
treating  the  contents  of  the  tube  with  hydrochloric  acid  and 
proving  that  hydrogen  is  being  evolved. 

Perform  the  same  experiments  using  a  mixture  of  copper 
oxide  and  finely  divided  carbon.  (See  page  191.) 

Experiment  118.  Preparation  of  Methane.  —  (Jena  glass  flask 
holding  from  100  to  150  cc. ;  glass  cylinders ;  trough ;  fused 
soda-lime;  fused  sodium  acetate.) 

Mix  20  grams  of  fused  soda-lime  with  20  grams  of  fused 
sodium  acetate,  and  powder  the  mixture  finely.  Both  sub- 
stances must  be  fused  to  remove  the  water,  which  would 
otherwise  be  set  free  during  the  experiment  and  break  the 
flask. 

Introduce  the  mixture  into  the  flask  and  connect  the  flask 
with  the  cylinder  as  shown  in  Figure  53.  Heat  the  flask 
strongly,  when  an  abundance  of  methane  will  be  given  off  and 
can  be  collected. 

When  a  flame  is  applied  to  the  mouth  of  a  cylinder  filled  with 
methane,  combustion  will  take  place.  (See  page  192.) 

Experiment  119.  Preparation  of  Acetylene.  —  (Calcium  car- 
bide.) 

A  few  small  pieces  of  calcium  carbide  are  thrown  into  a  dry 
test-tube,  and  a  few  drops  of  water  carefully  added.  Smell 
the  escaping  gas.  Touch  a  lighted  match  to  the  mouth  of  the 
test-tube,  and  the  acetylene  will  burn  with  a  smoky  flame.  (See 
page  193.) 

The  odor  of  acetylene  is  observed  when  a  bunsen  burner 
strikes  back  and  burns  at  the  bottom.  (See  page  193.) 

Experiment  120.  Carbon  Monoxide  prepared  by  Heating 
Oxalic  Acid  and  Sulphuric  Acid.  —  (Flask  holding  500  cc. ; 
thistle-tube ;  gas  wash- bottle  ;  glass  tubing ;  oxalic  acid ;  con- 
centrated sulphuric  acid;  sodium  hydroxide.) 

»  Arrange  the  apparatus  as  shown  in  Figure  54.  Into  the  flask 
Aj  20  grams  of  oxalic  acid  and  100  grams  of  concentrated  sul- 
phuric acid  are  introduced.  The  contents  of  the  flask  are  gently 
heated,  and  the  escaping  gases  —  carbon  monoxide  and  carbon 
dioxide — passed  through  the  wash-bottle  B  containing  a  solu- 
tion of  caustic  soda.  The  carbon  dioxide  is  absorbed  by  the 
alkali,  and  the  carbon  monoxide  passes  on.  After  the  air  has 
been  completely  driven  out  of  the  apparatus,  ignite  the  carbon 


212 


ELEMENTS  OF   INORGANIC   CHEMISTRY 


monoxide.  It  burns  with  a  slightly  luminous, 
bluish  flame.  What  is  formed  when  carbon  mo- 
noxide burns  ?  How  would  you  prove  it  ?  (See 
page  193.) 

Experiment  121.  Reducing  Power  of  Carbon 
Monoxide.  —  (Carbon  monoxide  generator  as  in 
the  last  experiment ;  hard-glass  tube ;  oxides  of 
several  of  the  heavy  metals,  such  as  iron,  copper, 
zinc,  lead,  etc.) 

Introduce  one  or  more 
of  these  oxides  at  differ- 
ent places  in  a  hard-glass 
tube  resting  on  a  combus- 
tion furnace ;  pass  the 
carbon  monoxide  through 
the  tube  and  heat  it  to  a 
high  temperature.  Are 
any  of  the  oxides  reduced 
to  the  corresponding 
metal  ?  How  do  you 
know  ?  (See  page  193.) 
Experiment  122.  Car- 
bon Dioxide  formed  by 
burning  Carbon  in  Oxygen.  —  (Gasometer  filled  with  oxygen ; 
gas  wash-bottle ;  hard-glass  tube ;  a  piece  of  charcoal ;  baryta 
water.) 

Introduce  a  piece  of  charcoal  into  a  hard-glass  tube  connected 
at  one  end  with  a  gasometer  filled  with  oxygen,  and  at  the 
other  with  a  gas  wash-bottle  containing  baryta  water. 

Pass  oxygen  through  the  tube  and  heat  the  charcoal.  It  will 
burn  vigorously  in  the  oxygen,  and  the  baryta  will  become 
cloudy  due  to  the  formation  of  a  precipitate  of  barium  carbon- 
ate, showing  that  carbon  dioxide  is  given  off.  (See  page  194.) 
Experiment  123.  Carbon  Dioxide  formed  when  Magnesite  is 
Heated.  —  (Hard-glass  tube,  closed  at  one  end ;  gas  wash-bottle ; 
baryta  water ;  magnesite  or  marble.) 

A  small  piece  of  magnesite  is  introduced  into  a  hard-glass 
tube,  connected  with- a  gas  wash-bottle  containing  baryta  water. 
The  magnesite  is  highly  heated,  when  carbon  dioxide  escapes 
and  forms  a  precipitate  of  barium  carbonate.  (See  page  195.) 
Experiment  124.  Carbon  Dioxide  in  the  Air;  in  large 
Quantities  in  Air  exhaled  from  the  Lungs.  —  (Gas  wash-bottle ; 
aspirator  as  in  Figure  34 ;  large  test-tube ;  baryta  water.) 


FIG.  54. 


CARBON,    SILICON,    BORON 


213 


By  means  of  an  aspirator  (Fig.  34)  draw  air  through  a  gas 
wash-bottle  containing  baryta  water.  Insoluble  barium  car- 
bonate will  gradually  be  formed. 

Close  a  large  test-tube  quickly  with  a  two-hole  rubber  stopper, 
after  introducing  a  dilute  solution  of  barium  hydroxide.     In- 
sert one  glass  tube  nearly  to  the  bottom  of  the  test-tube  as  in 
Figure   55.      A   glass 
tube      is     introduced  , — 
into  the  other  hole  in 
the  stopper.    Air  from 
the    lungs    is    driven 
through  the   solution. 
A  heavy  precipitate  of 
barium  carbonate  will 
quickly    form.       (See 
page  195.) 

Experiment  125. 
Carbon  Dioxide  formed 
by  the  Action  of  Acids 
on  Carbonates.  —  (So- 
dium carbonate;  cal- 
cium carbonate ;  dilute 
hydrochloric,  sulphu- 
ric, and  nitric  acids.) 

Treat   a  few  pieces 
of     marble     (calcium 
carbonate)    in   a  test- 
tube  with  hydrochloric  acid ;  vigorous  action  takes  place  and 
carbon  dioxide  escapes. 

Perform  the  same  experiment  using  nitric  acid.  Similar 
results  are  obtained.  Treat  sodium  carbonate  in  a  test-tube 
with  dilute  sulphuric  acid ;  carbon  dioxide  is  set  free. 

Carbon  dioxide  is  volatile,  and  is,  therefore,  driven  out  of  its 
salts  by  other  acids.  Furthermore,  it  is  a  very  weak  acid,  and 
for  this  reason  also  would  be  set  free  from  its  compounds  by 
other  acids.  (See  page  195.) 

Experiment  126.  Preparation  of  Carbon  Dioxide  on  the  Large 
Scale.  —  (Kipp  apparatus  ;  dilute  hydrochloric  acid ;  marble.) 

Carbon  dioxide  is  prepared  on  the  large  scale  by  the  action 
of  hydrochloric  acid  on  marble.  The  marble  is  introduced 
into  the  Kipp  apparatus  (Fig.  8),  and  also  dilute  hydro- 
chloric acid.  When  the  gas  is  desired,  it  is  only  necessary 
to  open  the  stop-cock.  When  no  more  gas  is  wanted,,  the 


FIG.  55. 


214  ELEMENTS  OF  INORGANIC  CHEMISTRY 

stop-cock  is  closed,  and  the  acid  driven  automatically  off  of 
the  marble. 

Experiment  127.  Carbon  Dioxide  will  not  burn  and  will  not 
support  Combustion.  —  An  attempt  to  ignite  carbon  dioxide 
will  result  in  failure.  A  lighted  taper  plunged  into  carbon 
dioxide  will  be  instantly  extinguished.  (See  page  197.) 

Experiment  128.  Carbon  Dioxide  Heavier  than  Air.  —  (Large 
balance  sensitively  adjusted  (Fig.  2) ;  two  large  beakers ; 
vessel  full  of  carbon  dioxide.) 

Attach  two  large  beakers  to  the  arms  of  the  large  balance 
(Fig.  2).  After  the  balance  has  been  brought  to  rest  with  the 
pointer  in  the  centre,  pour  the  carbon  dioxide  into  one  of  the 
beakers.  This  arm  of  the  balance  will  sink.  (See  page  197.) 

Experiment  129.  Solid  Carbon  Dioxide.  —  (Cylinder  con- 
taining liquid  carbon  dioxide ;  small  bag  of  thick  flannel ; 
small  porcelain  crucible  ;  cotton;  mercury.)  (This  experiment 
should  be  performed  only  by  the  instructor.) 

If  the  cylinder  containing  the  liquid  carbon  dioxide  is 
slanted  so  that  the  liquid  runs  to  the  front  of  the  tube,  and 
the  valve  is  carefully  opened,  some  of  the  carbon  dioxide  will 
be  volatilized,  while  a  part  will  be  solidified.  If  the  jet  is 
covered  with  a  thick  flannel  bag,  the  solid  carbon  dioxide  can 
be  caught,  and  we  have  carbon  dioxide  snow.  By  means  of 
this  substance  low  temperatures  can  be  produced.  Indeed, 
when  mixed  with  ether,  temperatures  lower  than  —80°  can 
be  produced. 

With  solid  carbon  dioxide  mercury  can  readily  be  frozen. 

Introduce  25  grams  of  mercury  into  a  small  porcelain  crucible, 
resting  on  a  poor  conductor  of  heat  such  as  cotton.  Cover  the 
mercury  with  solid  carbon  dioxide  and  cover  the  crucible.  In 
a  short  time  the  mercury  will  be  frozen. 

Remove  the  frozen  mercury  and  plunge  it  in  a  vessel  con- 
taining ice-water.  The  mercury  will  become  covered  with  a 
layer  of  ice. 

Experiment  130.  Alcohol  formed  by  the  Fermentation  of 
Sugar.  —  (Balloon  flask  holding  2  litres ;  two  glass  cylinders  ; 
condenser;  receiver;  grape  sugar;  yeast;  barium  hydroxide; 
sodium  hydroxide.) 

Arrange  the  apparatus  as  in  Figure  56.  Introduce  into  the 
flask  A  100  grams  of  grape  sugar  and  1  litre  of  water.  Add 
yeast,  and  set  in  a  warm  place.  Fermentation  will  soon  begin, 
and  carbon  dioxide  can  be  collected  in  cylinder  B,  which  con- 
tains baryta  water,  while  barium  carbonate  will  be  precipi- 


CARBON,   SILICON,   BORON 


215 


tated.    The  second  cylinder  C  contains  sodium  hydroxide,  to 
protect  the  baryta  from  the  carbon  dioxide  in  the  air. 


FIG.  56. 

After  the  flask  has  stood  for  several  days,  connect  with  a 
condenser  as  in  Figure  15,  and  distil  off  the  alcohol.  The  first 
portion  of  the  distillate  will  contain  a  mixture  of  alcohol  and 
water.  The  solution  will  probably  contain  too  little  alcohol  to 
take  fire  and  burn  when  a  flame  is  applied  to  it. 

Fill  a  test-tube  half  full  of  the  distillate,  and  add  potassium 
carbonate  and  shake  until  no  more  will  pass  into  solution. 
The  alcohol,  being  insoluble  in  a  saturated  solution  of  potassium 
carbonate,  will  rise  to  the  surface  and  appear  as  a  distinct  layer 
floating  on  the  solution.  It  can  be  removed  by  means  of  a 
pipette,  placed  in  a  small  evaporating-dish,  and  ignited.  (See 
page  198.) 

Experiment  131.  Carbon  Bisulphide,  Properties  of.  —  (Small 
evaporating-dish  ;  bellows ;  carbon  bisulphide.) 

That  carbon  bisulphide  is  very  inflammable  can  be  shown  by 
introducing  1  cc.  into  an  evaporating-dish  under  the  hood,  and 
bringing  near  the  dish  an  ignited  gas-lighter.  What  takes 
place  ? 

The  refrigerating  power  of  carbon  bisulphide  can  be  shown  by 
introducing  10  or  15  cc.  of  the  liquid  into  a  test-tube,  and  cover- 
ing this  with  1  or  2  cc.  of  water.  A  current  of  air  from  the 
blast-lamp  is  blown  through  the  mixture.  The  carbon  bisulphide 
evaporates  so  rapidly  that  the  water  is  frozen,  (See  page  198.) 


216  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Experiment  132.  Preparation  of  Cyanogen.  —  (Hard-glass 
tube  closed  at  one  end;  mercuric  cyanide.)  (TJiis  experiment 
should  be  performed  only  by  the  instructor.) 

Mercuric  cyanide  is  introduced  into  a  hard-glass  tube  closed 
at  one  end.  The  tube  is  drawn  out  to  a  fine  tube  at  the  other 
end.  Heat  the  mercuric  cyanide,  and  cyanogen  will  be  given 
off.  Ignite  the  jet  as  it  escapes  from  the  tube,  and  it  will  burn 
with  the  purple  flame  which  is  characteristic  of  the  substance. 
This  experiment  must  be  performed  only  under  a  good  hood, 
since  cyanogen  is  very  poisonous.  (See  page  198.) 

Experiment  1 33.  Flames  differ  greatly  in  their  Luminosity.  — 
(Candle ;  bunsen  burner ;  fish-tail  burner.) 

Light  a  candle,  a  bunsen  burner,  and  a  fish-tail  burner.  The 
candle  and  fish-tail  burner  are  brightly  luminous,  while  the 
bunsen  burner  gives  very  little  light.  The  fish-tale  burner  and 
the  bunsen  burner  both  use  gas,  and  yet  they  are  markedly 
different  in  their  light-giving  power. 

If  we  examine  these  burners  carefully,  we  shall  find  that  in  the 
bunsen  burner  the  gas  is  mixed  with  air  before  it  reaches  the 
flame,  and  thus  differs  from  the  fish-tail  burner.  The  combus- 
tion of  the  gas  is  more  perfect  in  the  bunsen  burner,  and  the 
illuminating  gas  is  diluted  with  an  indifferent  gas,  nitrogen, 
and  hence  the  lack  of  luminosity. 

If  the  air  is  cut  off  at  the  base  of  the  bunsen  burner,  the  flame 
at  once  becomes  luminous.  (See  page  200.) 

Experiment  134.  Rendering  a  Non-luminous  Flame  Lumi- 
nous. —  (Two  bunsen  burners ;  charcoal ;  platinum  cylinder.^) 

Ignite  a  bunsen  burner,  and  when  it  is  burning  with  a  practi- 
cally non-luminous  flame,  rub  together  two  pieces  of  carbon  near 
the  base  of  the  burner.  The  fine  carbon  dust  will  be  drawn 
into  the  burner  by  the  current  of  air,  and  the  flame  will 
become  strongly  luminous. 

Incline  a  bunsen  burner  at  an  angle  of  45  degrees,  and  place 
a  freshly  heated  platinum  tube  over  the  end  of  the  burner. 
Ignite  the  gas  at  the  end  of  the  platinum.  It  will  burn  with 
a  colorless  flame.  Now  heat  the  platinum  tube  to  redness  with 
a  second  burner,  and  the  first  flame  will  become  luminous.  (See 
page  201.) 

Experiment  135.  Use  of  the  Blowpipe.  —  (Mouth  blowpipe ; 
blowpipe  tube ;  a  piece  of  charcoal ;  several  heavy  metals  such 
as  zinc,  lead,  etc. ;  oxides  of  the  heavy  metals,  such  as  zinc 
oxide,  cadmium  oxide,  lead  oxide,  silver  oxide,  etc.) 

Bore  a  cavity  into  the  piece  of  charcoal  by  means  of  a  coin 


CARBON,    SILICON,   BORON  217 

placed  on  its  edge  and  rotated.  Place  a  small  piece  of  metal 
in  the  cavity,  arid  heat  with  the  oxidizing  flame  (the  outer  tip) 
of  the  blowpipe.  Describe  in  each  case  what  occurs  ?  Note 
any  incrustation  on  the  charcoal  ? 

In  a  similar  manner  place  the  oxide  of  a  heavy  metal  on  the 
charcoal,  and  heat  with  the  reducing  flame  of  the  blowpipe. 
Does  any  change  take  place  ?  Is  any  metal  formed  ?  Try  the 
action  of  acids  on  the  globule.  What  occurs?  (See  page  202.) 

Experiment  1'36.  Effect  of  cooling  the  Bunsen  Flame. —  (Bun- 
sen  burner ;  fine  wire  or  copper  gauze.) 

Ignite  a  bunseri  burner  and  lower  a  piece  of  wire  gauze 
upon  the  flame.  The  gas  will  burn  beneath  the  gauze  but  not 
above  it.  The  gauze  conducts  the  heat  away,  and  the  gas 
above  the  gauze  is  not  ignited.  After  a  time  the  gauze  be- 
comes sufficiently  hot  to  ignite  the  gas  above  it.  Apply  a 
match  to  the  gas  above  the  gauze ;  it  will  ignite. 

Turn  the  gas  on  a  bunsen  burner  and  do  not  ignite  the  gas. 
Bring  the  wire  gauze  about  5  cm.  above  the  burner,  and  ignite 
the  gas  above  the  gauze  with  a  match.  It  will  burn  above  the 
gauze  while  no  flame  will  appear  below  the  gauze.  (See 
page  203.) 

Experiment  137.  Preparation  of  Silicon  Dioxide.  —  (Water- 
glass  ;  hydrochloric  acid.) 

Treat  a  little  water-glass  with  hydrochloric  acid.  Filter  off 
the  precipitate,  wash,  dry,  and  ignite  in  a  porcelain  crucible. 
The  product  is  silicon  dioxide.  (See  page  205.) 

Experiment  138.  Formation  of  a  Salt  of  Silicic  Acid. — 
(Sand ;  porcelain  or  nickel  crucible  ;  sodium  carbonate.) 

Powder  a  little  sand  in  a  porcelain  mortar  until  it  is  a  mere 
dust.  Mix  it  with  ten  times  its  volume  of  sodium  carbonate, 
and  introduce  the  mixture  into  the  porcelain  or  nickel  crucible. 
Heat  over  a  blast-lamp  for  ten  minutes.  The  mass  will  then 
dissolve  in  water.  To  the  solution  of  sodium  silicate  add 
hydrochloric  acid.  What  is  precipitated  ?  (See  page  206.) 

Experiment  139.  Preparation  of  Boric  Acid;  Color  im- 
parted to  the  Alcohol  Flame.  —  (Borax ;  concentrated  hydro- 
chloric acid;  alcohol.) 

Dissolve  50  grams  of  borax  in  200  cc.  of  water,  and  add  con- 
centrated hydrochloric  acid  until  the  solution  is  strongly  acid  to 
litmus.  Wrhen  the  solution  cools,  boric  acid  crystallizes  out. 

Dissolve  a  few  crystals  of  the  boric  acid  in  a  little  alcohol, 
pour  the  solution  into  an  evaporating-dish  and  ignite  the 
alcohol.  It  will  burn  with  a  characteristic  greenish  flame. 


218  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Compare  this  flame  with  that  produced  by  burning  pure 
alcohol.     (See  page  208.) 


PROBLEMS 

1.  Amount  of  carbon  in  3  grams  of  methane?  in  3  grams 
of  ethylene  ?  in  3  grams  of  acetylene  ? 

2.  What  weight  of  carbon  dioxide  would  be  produced  by 
burning  a  diamond  weighing  0.5  grams  ? 

3.  Quantity  of  ethyl  alcohol  required  to  contain  1  gram  of 
carbon  ? 

4.  Weight  of  carbon  bisulphide  which  could  be  formed  from 
1  kilogram  of  graphite  and  sulphur ;  how  much  sulphur  would 
be  used  up? 

5.  When  3.3  grams  of  mercuric  cyanide  are  heated,  what 
would  be  the  weight  of  the  cyanogen  set  free  ? 

6.  Amount  of  silicon  in  7.3  grams  of  quartz  ? 

7.  Weight  of  boron  trioxide  in  100  grams  of  borax  ? 


CHAPTER  XVIII 

THE  METALS 

THE  metals  have  certain  properties  in  common  which  dis- 
tinguish them  from  the  other  elements.  With  one  exception, 
they  are  all  solids  at  ordinary  temperatures. 

Some  of  the  metals,  like  sodium,  potassium,  etc.,  are  very 
active  chemically,  while  others,  like  platinum,  gold,  etc.,  are 
very  resistant  to  chemical  reagents. 

The  chemistry  of  the  metals  is  in  general  much  simpler  than 
that  of  the  metalloids.  The  metals  form  ions  charged  with 
positive  electricity,  —  cations,  —  and  these  combine  with  the 
anions  of  acids,  forming  salts.  The  cations  are  generally  very 
much  simpler  than  the  anions,  consisting  usually  of  single 
metal  atoms  charged  with  positive  electricity. 

Since  when  metals  react  chemically  they  pass  into  solution, 
i.e.  into  the  ionic  state,  the  study  of  the  chemistry  of  the 
metals  is  largely  the  study  of  the  ions  which  they  form.  In- 
deed, we  have  excellent  reason  for  believing  that  in  order  that 
the  metals  should  react  chemically  they  must  be  in  the  ionic 
state.  If  this  be  true,  the  study  of  the  chemistry  of  the  metals  is 
in  reality  a  study  of  them  in  the  ionic  condition. 

We  shall  take  up  first  the  alkali  metals,  consisting  of 
lithium,  sodium,  potassium,  rubidium,  and  caesium. 

Next  in  order  come  the  metals  of  group  II.  These  fall, 
with  respect  to  their  relationships,  into  two  divisions,  —  cal- 
cium, strontium,  and  barium,  on  the  one  hand;  and  glucinum, 
or  beryllium  magnesium,  zinc,  cadmium,  and  mercury,  on  the 
other. 

When  we  pass  to  group  III,  we  find  that  boron  has  already 

219 


220  ELEMENTS   OF   INORGANIC   CHEMISTRY 

been  studied  with  the  metalloids.  The  first  metal  in  this 
group  is  aluminium.  In  the  same  group  are  the  rare  elements, 
scandium,  gallium,  yttrium,  indium,  lanthanum,  ytterbium, 
thallium,  and  samarium. 

Passing  to  iron,  we  have  in  this  same  group  nickel,  cobalt, 
manganese,  chromium,  and  the  rarer  elements,  molybdenum, 
tungsten,  and  uranium. 

Next  are  taken  up  copper,  silver,  and  gold,  and  then  lead 
and  tin ;  the  former  appearing  in  group  I,  the  latter  in 
group  IV. 

Finally,  among  the  noble  metals,  we  have  rhodium,  ru- 
thenium, palladium,  osmium,  iridium,  and  platinum. 

THE   ALKALI   METALS 

LITHIUM,   SODIUM,   POTASSIUM,   RUBIDIUM,   AND 
CESIUM 

SODIUM  (At.  Wt.  =  23.05) 

Occurrence  of  the  Element  Sodium.  —  The  element  sodium 
is  very  widely  distributed,  and  occurs  in  combination  with  other 
elements  in  many  places  in  large  quantities.  On  account  of 
its  great  chemical  activity  it  does  not  occur  in  nature  in  the 
free  condition.  Nearly  all  of  the  salts  of  sodium  are  soluble 
in  water.  We  should,  therefore,  expect  to  find  most  of  the 
sodium  compounds  dissolved  in  the  waters  of  the  sea,  and  such 
is  the  fact. 

In  certain  arid  regions  one  of  the  most  soluble  salts  of 
sodium  exists  in  large  beds.  In  Chili,  large  beds  of  sodium 
nitrate  are  found  which,  from  their  analogy  to  potassium 
nitrate  or  ordinary  saltpetre,  are  known  as  Chili  saltpetre. 

Sodium  salts  exist  in  great  abundance  in  certain  regions 
where  the  waters  of  the  sea  have  evaporated.  In  the  great 
salt  beds  of  the  earth,  such  as  those  at  Stassfurt,  the  chloride 
and  other  compounds  of  sodium  occur.  One  compound  of 
sodium  which  has  recently  come  into  prominence  in  connection 


THE   ALKALI  METALS  221 

with  the  manufacture  of  aluminium  should  be  mentioned. 
This  is  the  double  fluoride  of  sodium  and  aluminium,  Na3AlF6, 
occurring  in  Greenland  and  known  as  cryolite. 

Preparation  of  Sodium. —  Sodium  was  prepared  for  a  long 
time  by  the  reduction  of  the  oxide  or  hydroxide  by  means  of 
metallic  magnesium,  or  by  highly  heated  carbon. 

All  of  these  reduction  methods  are  now  abandoned  when  it 
is  desired  to  prepare  sodium  on  a  large  scale.  The  electrolytic 
method  is  used  entirely. 

Sodium  prepared  by  the  electrolysis  of  the  fused  hydroxide 
is  not  an  expensive  substance,  the  price  having  been  reduced 
immensely  by  the  application  of  the  electrolytic  process. 

Properties  of  Metallic  Sodium.  —  Sodium  is  a  soft  solid, 
which,  when  freshly  cut  with  a  knife,  has  a  metallic  lustre 
and  a  steel-gray  color.  The  surface  becomes  quickly  tarnished, 
due  to  the  rapidity  with  which  it  takes  up  oxygen  from  the 
air  or  from  moisture,  forming  the  oxide  or  hydroxide.  In  the 
presence  of  water  the  following  reaction  takes  place :  — 

2  Na  +  2  H.O  =  2NaOH  +  H2. 

The  hydrogen  which  is  liberated  when  a  piece  of  sodium  is 
thrown  upon  water,  does  not  take  fire  if  the  sodium  is  allowed 
to  move  about  over  the  surface  of  the  water.  If  the  sodium  is 
held  in  one  place,  as  by  throwing  it  upon  a  piece  of  filter- 
paper  upon  the  water,  enough  heat  is  produced  to  ignite  the 
hydrogen. 

Sodium  Hydride,  NaH,  is  formed  when  metallic  sodium  is 
heated  to  300°  in  a  current  of  hydrogen. 

Sodium  Peroxide,  NaO.  —  As  far  as  is  known  with  certainty, 
sodium  forms  only  one  compound  with  oxygen.  This  is  the 
peroxide  NaO.  It  is  obtained  when  sodium  is  heated  in  the 
atmosphere  to  about  300°  to  350°.  It  is  a  light  yellow  powder, 
and  dissolves  readily  in  water,  forming  sodium  hydroxide  and 
hydrogen  dioxide.  The  reaction  would  be  represented  thus  :  — 

2  NaO  +  2  H20  =  2  NaOH  +  H202. 


222  ELEMENTS  OF   INORGANIC   CHEMISTRY 

Sodium  Hydroxide,  NaOH. — We  have  just  seen  that  one 
method  of  preparing  sodium  hydroxide  is  to  treat  the  peroxide 
with  water.  Sodium  hydroxide  is  now  prepared  on  the  large 
scale  by  the  electrolysis  of  sodium  chloride.  Still  another 
method  is  to  treat  sodium  carbonate  with  calcium  hydrox- 

Na2C03  +  Ca(OH)2  =  CaC03  +  2NaOH. 

Another  method,  and  perhaps  the  best,  for  preparing  a  little 
sodium  hydroxide  in  very  pure  condition  is  to  allow  water  to 
act  on  metallic  sodium :  — 

2Na  +  2  H,0  =  2NaOH  +  H2. 

On  account  of  the  violence  of  this  reaction  it  must  be  carried 
out  with  certain  precautions. 

Sodium  hydroxide  is  one  of  the  strongest  bases  known. 
When  brought  into  the  presence  of  water,  it  dissociates 

thus :  —  +       _ 

NaOH  =  Na,  OH. 

When  sodium  hydroxide  is  treated  with  hydrochloric  acid, 
they  react  as  follows :  — 

Na,  0~H  +  H,  Cl  =  Na,  Cl  +  H2O. 

The  Chemistry  of  Sodium  the  Chemistry  of  the  Sodium  Ion.— 

The  chemistry  of  the  element  sodium  is  not  the  chemistry  of 
the  atom  or  molecule  of  sodium,  since  there  is  every  reason  for 
believing  that  these  are  practically  inert. 

On  the  other  hand,  wherever  we  have  sodium  ions  present, 
we  have  the  reactions  which  are  characteristic  of  this  element. 
Some  of  these  reactions,  together  with  their  products,  we  shall 
now  study. 

Sodium  Chloride  is  found  in  great  abundance  in  sea-water, 
which,  on  the  average,  contains  about  2.7  per  cent  of  the  salt. 
The  amount  of  sodium  chloride  in  sea-water,  however,  varies 
greatly  from  one  locality  to  another.  In  tropical  regions, 
where  the  evaporation  of  the  water  is  relatively  rapid,  the  per- 
centage may  be  as  much  as  3.5  to  3.8  per  cent. 


THE   ALKALI   METALS  223 

The  salt  is  obtained  from  sea-water  by  evaporation.  The 
water  is  allowed  to  flow  into  shallow  pools,  and  be  evaporated 
by  the  heat  of  the  sun. 

On  account  of  its  abundance  and  cheapness,  sodium  chloride 
is  of  great  importance  as  a  source  of  both  chlorine  and  sodium. 

Sodium  bromide,  NaBr,  and  sodium  iodide,  Nal,  resemble 
the  chloride  so  closely  in  their  properties  that  a  detailed  study 
of  them  in  this  connection  is  not  necessary. 

Sodium  Nitrate,  NaN03.  —  Sodium  nitrate  is  called  Chili 
saltpetre,  because  it  is  found  in  a  certain  rainless  district  in 
Chili,  and  since  it  is  the  sodium  analogue  of  potassium  nitrate 
or  ordinary  saltpetre.  It  is  extremely  soluble  in  water,  and, 
therefore,  could  not  exist  in  the  solid  condition  in  regions 
where  there  is  appreciable  rainfall. 

Sodium  nitrate  cannot  be  used  in  making  gunpowder  in  the 
place  of  potassium  nitrate,  since  it  absorbs  water  from  the  air, 
or  is  deliquescent,  as  we  say. 

Sodium  Sulphate,  NaJ504.10H20.  —  Sodium  sulphate,  called 
from  its  discoverer,  Glauber's  salt,  exists  in  certain  mineral 
waters  as  those  of  Carlsbad.  It  is  formed  in  a  large  number 
of  reactions.  When  sulphuric  acid  is  neutralized  with  sodium 
hydroxide,  we  have :  — 

2  NaOH  +  H2S04  =  2  H20  +  Na2S04. 

When  sodium  chloride  is  treated  with  sulphuric  acid,  sodium 
sulphate  is  formed,  not  because  sulphuric  acid  is  as  strong  as 
hydrochloric,  but  because  the  latter  is  volatile :  — 

2  NaCl  +  H2S04  =  2  HC1  +  JSTa^SO,. 

Acid  Sodium  Sulphate,  NaHS04. — Acid  sodium  sulphate  is 
formed  by  the  action  of  sulphuric  acid  on  salts  of  sodium  with 
volatile  acids :  — 

NaCl  +  H2S04  =  NaHS04  +  HC1. 

Also  by  the  action  of  sulphuric  acid  on  the  neutral  sulphate:  — 
Na2S04  +  H2S04  =  2  NaHS04. 


224  ELEMENTS   OF   INORGANIC    CHEMISTRY 

When  carefully  heated  in  a  vacuum  to  300°,  it  loses  water 
and  forms  the  pyrosulphate  :  — 

2  'NaHS04  =  H20  +  Na2S207. 

Sodium  Thiosulphate,  Na.,S203 .  5  H,0.  —  This  salt  is  frequently 
referred  to  as  sodium  hyposulphite,  or  in  commerce  simply  as 
"hypo."  It  is  prepared  by  dissolving  sulphur  in  a  solution  of 
sodium  sulphite :  — 

Na2S03+S=Na2S203. 

Solutions  of  this  salt  dissolve  silver  chloride  and  bromide, 
and  it  is,  therefore,  used  to  remove  these  substances  from  the 
photographic  plate  after  the  plate  has  been  exposed  to  the 
light. 

Sodium  Carbonate,  Na,C03 . 10H20.  — The  salt  with  ten  mole- 
cules of  water  crystallizes  from  a  solution  allowed  to  cool  on 
the  air.  This  salt  loses  a  part  of  its  water  at  ordinary  tem- 
peratures —  is  efflorescent. 

Tlie  Le  Blanc  Method  for  preparing  sodium  carbonate  was  used 
almost  exclusively  until  quite  recently.  In  the  Le  Blanc  method 
the  sodium  chloride  is  converted  into  the  sulphate  by  means  of 
sulphuric  acid.  The  sulphate  is  reduced  to  the  sulphide  by 
means  of  highly  heated  carbon,  the  sulphide  heated  with  cal- 
cium carbonate,  when,  at  a  sufficiently  high  temperature, 
calcium  sulphide  and  sodium  carbonate  are  formed.  The 
reactions  expressing  these  three  transformations  are:  — 

I.  2  NaCl  +  H2S04  =  2  HC1  +  Na2S04, 

II.  Na2S04  +  4  C  =  4  CO  +  Na2S, 

III.  Na2S  +  CaC03  =  CaS  +  Na2C03. 

It  is  not  difficult  to  separate  the  sodium  carbonate  from  the 
calcium  sulphide,  since  the  latter  is  difficultly  soluble  in  water, 
while  sodium  carbonate  is  readily  soluble. 

TJie  Solvay  or  Ammonia  Process  is  based  upon  the  fact  that 
acid  sodium  carbonate  is  much  less  soluble  in  water  than  acid 


THE   ALKALI   METALS  225 

ammonium  carbonate.  Ammonia  and  sodium  chloride  are  dis- 
solved in  water,  and  carbon  dioxide  passed  into  the  mixture. 
Under  these  conditions  acid  sodium  carbonate,  on  account  of 
its  small  solubility,  separates  from  the  solution.  The  reactions 
may  be  represented  thus  :  — 

NH3  +  H20  +  C02  =  NH4HC03, 
NH4HC03  +  NaCl  =  NH4C1  +  NaHC03. 

The  acid  sodium  carbonate  when  heated  forms  the  normal 
carbonate,  carbon  dioxide,  and  water :  — 

2  NaHC03  =  H20  +  C02  +  Na2C03. 

The  carbon  dioxide  is  conducted  into  more  ammonia  in  the 
presence  of  sodium  chloride,  and  the  process  is  thus  a  continu- 
ous one. 

Acid  Sodium  Carbonate,  NaHC03. — Primary  sodium  car- 
bonate, or  acid  sodium  carbonate,  or  the  "bicarbonate  of 
soda/'  is  formed  by  the  action  of  carbon  dioxide  on  the  normal 
carbonate :  — 

Na2C03  +  C02  +  H20  =  2  NaHC03. 

It  is  also  formed,  as  we  have  just  seen,  in  the  preparation  of 
normal  sodium  carbonate  by  the  ammonia  process.  When  acid 
ammonium  carbonate,  formed  by  the  action  of  carbon  dioxide 
on  ammonia,  is  treated  with  sodium  chloride,  acid  sodium  car- 
bonate is  formed :  — 

NH4HC03  +  NaCl  =  NaHC03  +  NH4C1. 

Hydrolysis  of  the  Carbonates.  —  The  aqueous  solution  of  nor- 
mal sodium  carbonate  has  a  strongly  alkaline  reaction.  Indeed, 
the  primary  or  acid  carbonate  has  a  weakly  alkaline  reaction. 
This  is  due  to  the  presence  of  hydroxyl  ions  in  the  aqueous 
solutions  of  the  carbonates.  The  carbonates  are  salts  of  the 
very  weak  carbonic  acid,  and  like  all  salts  of  weak  acids  are 
hydrotyzed  to  a  greater  or  less  extent  by  water :  — 

Ka2C03  +  H20  =  Na,  OH  +  Na,  HC03. 


226  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Even  the  acid  carbonate  is  hydrolyzed  to  a  sufficient  extent  to 
show  an  alkaline  reaction. 

All  carbonates  which  are  soluble  in  water  are  hydrolyzed 
sufficiently  to  show  an  alkaline  reaction. 

The  Phosphates  of  Sodium.  —  Since  phosphoric  acid  is  tribasic 
(see  phosphorus),  there  are  three  sodium  salts  of  this  acid  pos- 
sible, and  all  are  known.  The  secondary  sodium  phosphate, 
Na2HP04,  is  by  far  the  best  known,  and  is  always  meant  when 
the  term  "  sodium  phosphate  "  is  used  without  qualification.  It 
contains  twelve  molecules  of  water  of  crystallization. 

Its  aqueous  solution  is  slightly  alkaline,  due  to  the  fact  that 
phosphoric  acid  is  a  weak  acid,  and  it  is  slightly  hydrolyzed  by 
water  :  —  +  _  + 

Na2HP04  +  H20  =  Na,  OH  +  Na,  H2P04. 

When  disodium  phosphate  is  treated  with  one  equivalent 
of  phosphoric  acid,  the  monosodium  phosphate,  NaH2P04,  is 

Na2HP04  +  H3P04  =  2  NaH2P04. 

If,  on  the  other  hand,  one  equivalent  of  sodium  hydroxide 
is  added  to  one  equivalent  of  disodium  phosphate,  the  trisodium 
phosphate  is  formed  :  — 

Na2HP04  4-  NaOH  =  Na3P04  +  H20. 


Sodium  Borate  or  Tetraborate,  Na-JJA.  10  H20.  —  Borax  is  not 
the  salt  of  normal  boric  acid,  H3B03,  but  of  a  poly  boric  acid, 
derived  from  the  normal  acid  by  loss  of  water.  When  four 
molecules  of  boric  acid  lose  five  molecules  of  water,  the  acid 
from  which  borax  is  formed  results  :  — 

4H3B03=5H20  +  H2B407. 

Borax,  or  sodium  tetraborate,  is  formed  when  boric  acid  is 
neutralized  with  sodium  carbonate  :  — 


4  H3B03  +  Na2C03  =  Na«B407  +  6  H20 

On  account  of  the  property  of  dissolving  metal  oxides  borax 
is  frequently  used  as  &flux.     It  is  used  for  the  same  reason  to 


THE  ALKALI   METALS  227 

clean  two  metal  surfaces  which  it  is  desired  to  solder  together. 
These  usually  become  covered  with  a  layer  of  oxide  when  the 
metal  is  heated,  and  the  solder  will  not  adhere  to  the  surfaces 
while  the  oxide  is  present.  When  a  little  borax  is  added,  it 
removes  at  the  elevated  temperature  the  oxides  already  formed, 
and  at  the  same  time  protects  the  hot  metal  surfaces  from  the 
oxygen  of  the  air. 

POTASSIUM  (At.  Wt.  =  39.14) 

Occurrence  and  Preparation.  —  Potassium  like  sodium  does 
not  occur  in  nature  in  the  free  condition,  and  for  the  same 
reason,  viz.  its  great  chemical  activity. 

Plants  have  the  power  of  taking  up  potassium  ions  in  large 
quantities  and  building  them  up  into  their  tissues.  When  such 
plants  are  burned  the  potassium  salts  remain  behind  in  the 
ashes. 

Potassium  also  occurs  in  many  of  the  more  common  rocks  and 
minerals  in  the  form  of  silicates.  Ordinary  feldspar  is  a  double 
silicate  of  potassium  and  aluminium. 

Potassium  salts  also  occur  in  the  great  salt  beds,  especially 
in  those  of  Stassfurt,  in  Germany. 

Potassium  was  prepared  by  reducing  the  sulphide  or  hydrox- 
ide with  highly  heated  metals,  such  as  magnesium,  aluminium, 
iron,  etc. 

All  of  these  methods  have  now  been  abandoned  in  favor 
of  the  electrolytic.  Metallic  potassium  is  now  prepared  by 
electrolyzing  the  chloride,  or  hydroxide. 

Properties  of  Potassium.  — Potassium  is  characterized  by  its 
great  chemical  activity,  being  even  more  active  than  sodium. 
When  a  small  piece  of  potassium  is  thrown  upon  water,  it 
decomposes  it  in  the  same  manner  as  sodium,  yielding  potas- 
sium hydroxide  and  setting  hydrogen  free :  — 

2  H20  +  2  K  =  2  KOH  +  H2. 

The  action  in  the  case  of  potassium  is,  however,  so  vigorous 
that  even  when  the  metal  is  allowed  to  move  around  over  the 


228  ELEMENTS  OF  INORGANIC  CHEMISTRY 

surface  of  the  water,  enough,  heat  is  generated  to  ignite  the 
hydrogen. 

Potassium  Hydride,  KH,  is  formed  when  potassium  is  heated 
to  300°  —  400°  in  a  current  of  hydrogen. 

Potassium  Peroxide,  K02.  —  The  only  compound  which  potas- 
sium is  known  to  form  with  oxygen  is  the  dioxide.  It  is 
obtained  by  heating  potassium  in  a  current  of  dry  oxygen.  It 
is  an  orange-colored  powder,  melting  at  280°.  In  contact  with 
water  it  forms  potassium,  hydroxide,  hydrogen  dioxide,  and 
oxygen : — 

2  K02  4-  2  H20  =  2  KOH  +  H202  +  02. 

Potassium  Hydroxide,  KOH.  —  When  metallic  potassium  is 
thrown  upon  water,  potassium  hydroxide  is  formed,  as  we  have 
seen. 

It  is  also  formed  when  the  peroxide  is  dissolved  in  water. 

When  a  concentrated,  aqueous  solution  of  potassium  chloride 
is  electrolyzed,  potassium  hydroxide  is  formed  in  the  solution. 

Potassium  hydroxide  dissolves  very  readily  in  water,  forming 
caustic  potash,  and  the  solution  is  one  of  the  strongest  bases 
known.  It  dissociates  completely  into  potassium  and  hydroxyl 

ions,  —  + 

KOH  =  K,  OH, 

and  this  at  no  very  great  dilution.     It  is  also  formed  by  the 
action  of  lime-water  on  potassium  carbonate :  — 

K2C03  +  Ca(OH)2  =  CaC03  +  2  KOH. 

Potassium  hydroxide,  on  account  of  its  great  solubility, 
readily  precipitates  the  hydroxides  of  the  heavy  metals  from 
aqueous  solutions  of  their  salts  :  — 

A+gN~03  +  K,  OH  =  AgOH  +  K,  N03. 

Silver  hydroxide,  however,  breaks  down  into  silver  oxide  and 
water. 

Potassium  Chloride,  KC1,  occurs  in  nature  as  such,  in  com- 
bination with  magnesium,  chloride  as  the  mineral  carnallite, 
KMgCl3 . 6  H20,  and  in  other  combinations.  When  a  hot 


THE  ALKALI  METALS  229 

solution  of  this  salt  crystallizes,  the  double  salt  decomposes, 
potassium  chloride  separating  out.  It  is  a  beautifully  white 
substance,  crystallizing  in  cubes,  which  are  readily  dissolved 
by  water.  Potassium  chloride  is  a  type  of  a  salt  of  a  strong 
acid  with  a  strong  base.  We  have  seen  that  strong  acids  and 
strong  bases  mean  those  that  are  greatly  dissociated.  The 

compound  formed  by  the  union  of  the  cation  of  the  base,  K, 
with  the  anion  of  the  acid,  Cl,  is  among  the  most  strongly  dis- 
sociated substances  known.  A  dilute  solution  of  potassium 
chloride  is,  therefore,  a  solution  of  potassium  and  chlorine  ions 
and  nothing  else,  there  being  no  molecules  in  the  solution. 
All  the  properties  of  such  solutions,  both  chemical  and  physi- 
cal, are  the  properties  of  chlorine  ions  and  potassium  ions, 
since  these  only  are  present.  The  bromide,  iodide,  and  fluoride 
of  potassium  resemble  closely  the  chloride. 

Potassium  Chlorate,  KC103  —  Potassium  combines  with  the 
oxygen  acids  of  chlorine,  forming  well-defined  salts.  A  few  of 
these  are  of  sufficient  importance  to  merit  special  considera- 
tion. Potassium  chlorate  is  prepared,  as  we  have  seen,  by  the 
action  of  chlorine  on  caustic  potash :  — 

6  KOH  +  3  C12  =  5  KC1  +  KC103  +  3  H20. 

It  is  separated  from  the  chloride  by  its  solubility  in  water 
being  much  less  than  that  of  the  chloride. 

When  potassium  chlorate  is  heated,  it  decomposes  in  the 
sense  of  the  following  equation, — 

2  KClOg  =  KC1  +  KC104  +  02, 

potassium  chloride  and  perchlorate  being  formed.  When  the 
perchlorate  is  heated  to  a  still  higher  temperature,  it  breaks 
down  into  the  chloride  and  oxygen. 

When  potassium  chlorate  is  powdered  with  a  small  piece  of 
sulphur,  an  explosion  occurs,  which  is  violent  if  an  appreciable 
quantity  of  sulphur  is  used.  A  violent  explosion  results  if 
potassium  chlorate  is  brought  together  with  phosphorus.  With 
antimony  sulphide  an  explosive  mixture  is  also  formed. 


230  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Potassium  chlorate  is  extensively  used  in  the  preparation  of 
matches.  The  so-called  safety  matches  are  made  of  a  mixture 
of  potassium  chlorate  and  sulphide  of  antimony.  When  these 
are  rubbed  upon  a  surface  covered  with  red  phosphorus,  a 
miniature  explosion  results  and  the  whole  mass  is  ignited. 
When  rubbed  upon  an  ordinary  object  such  matches  do  not 
take  fire. 

Potassium  Nitrate,  KN03.  —  Potassium  nitrate  or  saltpetre  is 
one  of  the  most  important  salts  of  potassium.  It  is  very 
soluble  in  water  and,  therefore,  does  not  occur  in  the  solid 
state  in  any  considerable  quantity  in  regions  where  there  is  an 
abundant  rainfall. 

Potassium  nitrate  is  formed  in  large  quantity  in  the  saltpetre 
plantations.  Refuse  animal  matter  which  contains  nitrogen  is 
mixed  with  potassium  carbonate,  and  exposed  to  the  action 
of  the  nitrifying  ferment  or  saltpetre  bacteria  in  the  soil. 
The  oxygen  of  the  air,  through  the  agency  of  these  bacteria, 
oxidizes  the  ammonia  formed  from  the  decomposing  organic 
matter,  to  nitric  acid,  which  then  combines  with  potassium 
hydroxide  or  carbonate  and  forms  the  corresponding  nitrate. 

Potassium  nitrate  is  made  to-day  chiefly  from  sodium  nitrate 
or  Chili  saltpetre.  When  a  solution  of  sodium  nitrate  is  mixed 
with  a  solution  of  potassium  chloride,  the  following  reaction 
takes  place :  — 

Na,  N03  +  K,  Cl  =  NaCl  +  K,  NO,. 

The  sodium  chloride  is  deposited  at  higher  temperatures. 
When  the  solution  is  allowed  to  cool  down,  potassium  nitrate 
is  deposited.  Gunpowder  is  a  mixture  of  potassium  nitrate, 
sulphur,  and  carbon,  in  such  proportions  as  to  secure  complete 
combustion.  The  nitrate  gives  off  oxygen,  which  combines 
with  the  carbon,  forming  carbon  dioxide  ;  the  nitrogen  escapes 
as  such,  and  the  potassium  remains  behind  in  the  form  of  the 
sulphide  or  sulphate.  The  equation  which  is  usually  written  to 
express  the  decomposition  of  gunpowder  is  :  — 

2  KN03  +  3  C  +  S  =  K2S  +  3  C02  +  N2. 


THE  ALKALI  METALS  231 

Gunpowder  is  prepared  by  mixing  the  three  constituents  in 
the  following  proportions :  — 

KN03 75  per  cent 

S 12  per  cent 

C 13  per  cent 

This  corresponds  almost  exactly  to  three  molecules  of  salt- 
petre, three  atoms  of  carbon,  and  one  atom  of  sulphur ;  and  is 
the  chief  reason  for  writing  the  above  very  simple  equation. 

When  gunpowder  is  ignited,  the  gases  liberated  occupy  several 
hundred  times  the  volume  of  the  powder ;  or  if  they  are  forced 
to  occupy  the  same  volume  as  the  original  powder,  the  pressure 
exerted  is  several  hundred  atmospheres.  This  is  the  principle 
made  use  of  in  employing  explosives  to  drive  missiles  with  a 
high  velocity.  The  gunpowder  is  exploded  in  a  metal  tube 
closed  on  all  sides  and  open  at  one  end.  The.  ball  is  placed 
tightly  upon  the  powder,  so  that  when  the  latter  explodes  the 
gases  are  liberated  in  practically  a  closed  space.  An  enormous 
pressure  is  thus  generated,  which  drives  the  ball  out  of  the  end 
of  the  gun  with  a  high  velocity. 

Potassium  Sulphate,  K2S04,  occurs  in  nature  in  combination 
with  magnesium  sulphate  and  magnesium  chloride  as  kainite. 
It  occurs  in  a  number  of  the  salt  beds,  but  especially  in  those 
at  Stassfurt  and  other  places  in  Germany.  Kainite  has  the  com- 
position K2S04 .  MgS04 .  MgCl2 .  6  H2O.  The  double  sulphate  of 
potassium  and  magnesium  is  treated  with  chloride  of  potassium, 
when  magnesium  chloride  and  potassium  sulphate  result. 

When  normal  potassium  sulphate  is  treated  with  an  equivalent 
of  sulphuric  acid,  the  acid,  or  primary  sulphate,  is  formed :  — 

K2S04  +  H2S04  =  2  KHS04. 

Acid  potassium  sulphate  in  aqueous  solution  shows  a  strongly 
add  reaction.  This  is  due  to  the  fact  that  sulphuric  acid  is  a  strong 
acid,  and  the  hydrogen  of  the  acid  sulphate  begins  to  dissociate  : 

4  =  K,H,S04. 


232  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Potassium  Carbonate,  K2C03.  —  Potassium  carbonate  was  ob- 
tained for  a  long  time  mainly  from  the  ashes  of  certain  plants. 
When  the  plants  were  burned,  the  potassium  remained  behind 
in  the  form  of  the  carbonate.  This  was  obtained  by  leaching 
the  ashes  with  water  and  evaporating  the  solution,  when  the 
impure  carbonate  crystallized  out.  Potassium  carbonate  is 
also  obtained  from  the  residues  of  the  beet-sugar  industry. 

The  carbonate  is  purified  by  means  of  the  difference  in  solu- 
bility between  the  impurities  and  the  salt  in  question. 

Acid,  or  Primary  Potassium  Carbonate,  KHC03.  —  The  acid 
salt  is  formed  by  conducting  carbon  dioxide  into  the  solution 
of  the  neutral  salt :  — 

K2C03  +~C02  +  H20  =  2  KHC03. 

Dilute  solutions  of  acid  potassium  carbonate  react  alkaline. 
This  is  due  to  the  hydrolysis  of  the  acid  salt  by  water,  forming 
hydroxyl  ions :  — 

KHC03  +  H20  =  K,  OH  +  H2C03. 

Phosphates  of  Potassium.  —  Potassium,  like  sodium,  forms 
three  salts  with  phosphoric  acid,  —  the  primary,  KH2P04; 
secondary,  K2HP04,  and  tertiary,  K3P04,  phosphates. 

These  phosphates  do  not  call  for  any  special  comment. 

The  elements  lithium,  rubidium,  and  ccesium  are  compara- 
tively rare  elements  which  need  not  be  considered  in  detail. 

AMMONIUM 

The  group  ammonium,  although  not  an  element,  closely 
resembles  in  its  properties  the  alkali  metals.  It  forms  a  univ- 
alent  cation,  NH4,  which  has  the  power  to  combine  with  the 
anions  of  acids  and  form  salts  which  resemble  in  many  respects 
those  of  the  alkali  metals. 

Ammonium  Hydroxide,  NH4OH.  —  Ammonia  combines  with 
water,  forming  the  hydroxide  NH4OH  :  — 


THE   ALKALI   METALS  233 

In  the  presence  of  water  this  compound  is  dissociated  to  some 
extent  into  the  ammonium  ion,  NH4,  and  the  hydroxyl  ion, 

NH4OH  =  NH4,  OH. 

It  is,  therefore,  a  base,  but  it  is  a  very  weak  base. 

When  an  aqueous  solution  of  ammonium  hydroxide  is  boiled, 
it  breaks  down  into  ammonia  and  water  :  — 


This  fact  is  made  use  of  in  detecting  the  presence  of  ammonia 
or  an  ammonium  salt.  The  ammonium  salt  is  treated  with  a 
strong  base  like  caustic  soda,  when  it  is  broken  down  into  the 
sodium  salt,  and  ammonia  which  is  given  off  when  the  solution 
is  heated.  This  can  be  detected  by  the  odor  when  present  in 
considerable  quantity,  or  by  holding  a  piece  of  moistened  red 
litmus  in  the  escaping  vapors  when  the  ammonia  is  present  in 
small  quantity.  This  becomes  colored  blue. 

Ammonium  Chloride,  NH4C1.  —  When  hydrochloric  acid  is 
neutralized  with  ammonium  hydroxide  and  the  solution  evapo- 
rated, ammonium  chloride  or  sal  ammoniac  is  obtained.  This 
salt  is  a  beautifully  crystalline  compound,  which  is  readily 
soluble  in  water.  Although  ammonium  hydroxide  is  only 
slightly  dissociated,  the  salt  with  hydrochloric  acid  is  among 
the  most  strongly  dissociated  compounds.  This  is  true  in 
general  of  the  salts  of  ammonia  with  strong  acids. 

Ammonium  Hydrosulphide,  Sulphide,  and  Polysulphides.  — 
When  a  solution  of  ammonium  hydroxide  is  saturated  with 
hydrogen  sulphide,  the  hydrosulphide,  NH4HS,  is  produced:  — 

NH4OH  +  H2S  =  H20  +  NH4HS. 

The  sulphide  of  ammonium,  (NH4)2S,  is  formed  by  treating  the 
hydrosulphide  in  solution  with  an  equivalent  of  ammonia:  — 

NH4HS  +  NH4OH  =  (NH4)2S  +  H20. 

The  aqueous  solution  of  ammonium  sulphide  is  colorless 
when  freshly  prepared,  but  when  allowed  to  stand  for  a  time 


234  ELEMENTS  OF  INORGANIC   CHEMISTRY 

it  becomes  deep  yellow  in  color.  This  is  due  to  the  oxidation 
of  the  sulphide  by  the  oxygen  of  the  air  setting  sulphur  free :  — 

(NH4)2S  +  0  =  H20  +  2  NH3  +  S. 

The  sulphur  thus  set  free  acts  on  more  ammonium  sulphide, 
forming  the  polysulphides  of  ammonium.  There  are  several 
of  these  compounds. 

Ammonium  Sulphate,  (NH4)2S04.  —  This  is  one  of  the  most 
important  salts  of  ammonia,  as  being  one  of  the  chief  sources 
of  the  ammonia  used  on  a  commercial  scale. 

Ammonium  Carbonate,  (NH4)2C03 .  H20. — When  a  mixture 
of  calcium  carbonate  and  ammonium  sulphate  is  distilled,  and 
ammonia  passed  into  the  aqueous  solution  of  the  product,  nor- 
mal ammonium  carbonate  is  formed.  This  is  not  very  stable, 
and  breaks  down  readily  into  ammonia  and  the  acid  carbonate, 

NH4HC03:  — 

(NH4)2C03  =  NH3  +  NH4HC03. 

The  acid  carbonate  is  also  formed  by  the  action  of  carbon  di- 
oxide on  aqueous  ammonia.  This  is  a  much  more  stable  sub- 
stance than  the  normal  carbonate. 

The  Double  Phosphate  of  Ammonium  and  Sodium.  —  The 
double  phosphate  of  ammonium  and  sodium,  NaNH4HP04,  or 
microcosmic  salt,  is  formed  by  bringing  together,  in  solution, 
disodium  phosphate  and  ammonium  chloride :  — 

Na2HP04  +  NH4C1  =  NaCl  +  NH4NaHP04. 
When  heated  it  decomposes  into  ammonia,  water,  and  sodium 
metaphosphate :  — 

NH4NaHP04  =  NH3  +  H20  +  KaP03. 

Characteristics  of  the  Alkali  Metals  in  General.  —  From  the 
foregoing  study  of  the  alkalies  we  can  draw  general  conclusions 
as  to  their  chemical  behavior.  In  the  first  place,  they  are  all 
strongly  base-forming  elements,  which  is  the  same  as  to  say  that 
when  they  are  dissolved  in  water  they  form  strongly  electro- 
positive cations,  there  being  a  corresponding  number  of  hy- 
droxyl  anions  present  in  the  solution. 


THE   ALKALI  METALS  235 

The  alkalies  form  only  univalent  ions,  which  means  that  they 
can  carry  only  one  electrical  charge. 


EXPERIMENTS  WITH  SODIUM  POTASSIUM  AND 
AMMONIUM 

Experiment  140.  Properties  of  Sodium.  —  (Small  piece  of 
metallic  sodium.) 

Eemove  with  forceps  a  piece  of  metallic  sodium  the  size  of 
a  pea  from  the  bottle  containing  sodium  covered  with  kerosene. 
Dry  it  by  pressing  between  layers  of  filter-paper.  Do  not  touch 
metallic  sodium  with  the  hand. 

Holding  the  sodium  in  a  pair  of  forceps,  cut  it  with  a  dry 
knife.  What  is  the  appearance  of  the  surface?  Does  it 
change  ?  Why  ? 

Throw  a  piece  of  sodium  the  size  of  a  grain  of  rice  on  water 
in  an  evapprating-dish.  What  do  you  observe  ? 

Lay  a  piece  of  filter-paper  on  the  water,  and  when  it  is  wet 
throw  a  piece  of  sodium  on  the  paper.  The  hydrogen  takes  fire. 
Why  ?  (See  page  221.) 

Experiment  141 .  Properties  of  Sodium  Hydroxide.  —  (Sodium 
hydroxide.) 

Dissolve  sodium  hydroxide  in  water  in  which  a  thermometer 
is  immersed.  Is  there  any  change  in  the  temperature  ?  Test 
the  reaction  towards  red  litmus  ?  What  is  formed  when  sodium 
hydroxide  is  neutralized  with  acids  ?  (See  page  222.) 

Experiment  142.  Preparation  of  Sodium  Carbonate  by  the 
Ammonia  or  Solvay  Process.  —  (Ammonia ;  sodium  chloride ; 
carbon  dioxide  generator.) 

Make  a  fairly  concentrated  solution  of  sodium  chloride  in 
100  cc.  of  aqueous  ammonia.  Pass  carbon  dioxide  into  the 
solution  until  no  more  is  absorbed.  A  white  precipitate  is 
formed.  Dry  the  compound  and  then  place  a  little  in  a  hard- 
glass  tube  closed  at  one  end,  and  connected  with  a  bottle  con- 
taining baryta  water.  Heat  the  compound.  Is  carbon  dioxide 
given  off  ? 

When  the  gas  ceases  to  escape,  remove  the  residue  and 
treat  it  with  an  acid.  Is  carbon  dioxide  given  off?  Ex- 
plain why  all  the  carbon  dioxide  was  not  driven  off  by  heat. 
Write  the  equations  for  all  of  the  above  reactions  ?  (See 
page  225.) 


236 


ELEMENTS   OF   INORGANIC    CHEMISTRY 


Experiment  143.  Potassium  Hydroxide  precipitates  the 
Hydroxide  of  the  Heavy  Metals.  —  (Ferric  chloride;  lead  ni- 
trate ;  nickel  chloride  ;  copper  chloride.) 

Make  dilute  solutions  of  the  above-named  salts  in  water, 
and  add  a  few  drops  of  potassium  hydroxide  to  each  solution. 
Describe  in  each  case  what  you  see  ? 

Sodium  hydroxide  will  answer  just  as  well  as  potassium 
hydroxide.  (See  page  228.) 

Experiment  144.  Quantitative  Determination  of  Oxygen 
in  Potassium  Chlorate.  —  (Hard-glass  tube  8  mm.  in  diameter, 
closed  at  one  end  ;  burette,  Figure  58  ;  potassium  chlorate.) 

Weigh  a  hard-glass  tube  of  about  8  mm.  internal 
diameter,  and  12  cm.  long,  closed  at  one  end.  Intro- 
duce about  0.2  gram  of  potassium  chlorate,  and  re  weigh 
the  tube  and  chlorate  so  as  to  ascertain  the  exact 
weight  of  the  chlorate  in  the  tube.  Introduce  a  plug 
of  freshly  heated  asbestos  and  shove  it  down  upon 
the  chlorate  in  the  tube.  This  will  sweep  down  any 
chlorate  adhering  to  the  walls  of  the  tube. 

Draw  the  tube  out  as  shown  in  the 
figure  (57). 

Attach  to  a  burette  of  the  form 
shown  in  Figure  58,  after  filling  the 
tubes  nearly  full  of  water.  Bring  the 
water  to  the  same  level  in  the  two 
arms  of  the  burette  and  read  the 
burette. 

Now  heat  the  potassium  chlorate 
until  all  of  the  oxygen  has  been 
driven  off.  Allow  the  gas  to  stand 
in  the  burette  until  it  has  come  to 
^3  same  temperature  as  the  air  in 
the  room.  Bead  the  burette  again  after  bring- 
ing the  water  in  both  arms  to  the  same  level. 
Also  read  a  thermometer  suspended  from  the 
top  of  the  burette,  and  a  barometer,  in  the 
same  room. 

Correct  the   observed  volume  of  the   gas 
according  to  the  equation,  page  112. 
Knowing  the  weight  of  a  litre  of 
oxygen  (1.4296  grams),  calculate  the 
weight  of  the  oxygen  obtained. 

Calculate  the  weight  of  the  oxy-    jj$r  FIG,  58, 


FIG.  57. 


THE  ALKALI  METALS  237 

gen  which  should  be  given  off  by  the  potassium  chlorate 
used.  Compare  the  two  results.  (See  page  229.) 

Experiment  145.  Potassium  Nitrate  mixed  with  Carbon 
readily  burns.  —  (Charcoal ;  potassium  nitrate.) 

Powder  3  grams  of  charcoal  in  a  mortar,  and  mix  intimately 
with  it  18  grams  of  powdered  potassium  nitrate.  Place  the 
mixture  in  a  metal  pan  and  ignite  it  with  the  flame  of  a  bunsen 
burner.  What  takes  place  ?  (See  page  230.) 

Experiment  146.  Volatility  of  Ammonium  Salts.  —  (Evapo- 
rating-dish;  porcelain  triangle;  ammonium  chloride,  sul- 
phate, nitrate,  and  carbonate.) 

Place  a  small  porcelain  evaporating-dish  on  a  porcelain  tri- 
angle on  a  tripod,  and  introduce  one  of  the  above  salts.  Heat, 
and  it  completely  volatilizes.  Perform  the  same  experiment 
with  the  other  salts.  (See  page  233.) 

Experiment  147.  Formation  of  Ammonium  Sulphide  and 
Ammonium  Hydrosulphide.  —  (Ammonia;  hydrogen  sulphide.) 

Divide  100  cc.  of  aqueous  ammonia  into  two  equal  parts. 
Saturate  one-half  with  hydrogen  sulphide.  What  is  formed  ? 
Add  the  other  half.  What  is  formed  ?  {See  page  233.) 

Experiment  148.  Ammonium  Sulphide  precipitates  the  Sul- 
phides of  the  Heavy  Metals.  —  (Solution  of  ammonium  sul- 
phide; solution  of  lead  nitrate,  cadmium  chloride,  zinc  chlo- 
ride, mercuric  chloride,  silver  nitrate.) 

Add  ammonium  sulphide  to  solutions  of  each  of  the  above 
substances.  Describe  in  each  case  what  you  observe?  (See 
page  233.) 

Experiment  149.  General  Flame  Reactions  of  Sodium  and 
Potassium.  —  (Salts  of  sodium ;  salts  of  potassium ;  ammonium 
salts;  hydrofluosilicic  acid;  perchloric  acid;  hydrochlorpla- 
tinic  acid ;  tartaric  acid ;  4  platinum  wires  each  10  cm.  long.) 

Dip  a  platinum  wire  into  hydrochloric  acid  and  then  heat  in 
a  bunsen  flame  until  it  no  longer  imparts  color  to  the  flame. 
Then  dip  the  wire  into  any  sodium  salt  and  thrust  it  into  the 
flame.  How  does  sodium  affect  the  flame  ? 

Dip  another  clean  platinum  wire  into  any  potassium  salt  and 
thrust  it  into  the  flame.  What  color  does  potassium  impart 
to  the  flame  ? 

Mix  a  potassium  salt  with  a  sodium  salt  and  dip  a  clean 
platinum  wire  into  the  mixture.  The  yellow  sodium  light 
being  so  intense,  the  purple  due  to  the  potassium  may  not  be 
seen.  Hold  between  the  eye  and  the  flame  a  piece  of  blue 


238  ELEMENTS   OF  INORGANIC   CHEMISTRY 

cobalt  glass.  This  will  cut  off  the  yellow  light  due  to  the 
sodium,  and  then  the  purple  potassium  flame  can  be  readily 
seen.  This  is  a  convenient  means  of  distinguishing  between 
sodium  and  potassium,  and  of  detecting  each  in  the  presence 
of  the  other  when  both  are  present. 

Ammonium  salts  do  not  give  any  characteristic  color  to  the 
flame. 

PROBLEMS 

1.  Weight  and  volume  of  hydrogen  that  would  be  liber- 
ated when  2  grams  of  sodium  act  on  water  ?     What  amount  of 
sodium  chloride  could  be  produced  from  the  sodium  hydroxide 
formed  ? 

2.  How  much  sodium  sulphate  would  be  formed  when  500 
grams  of  sodium  chloride  are  treated  with  sulphuric  acid  ? 

3.  Percentage  of  potassium,  oxygen,  and  nitrogen  in  potas- 
sium nitrate  ? 

4.  From  7.3  grams  of  microcosmic  salt  how  much  sodium 
metaphosphate  can  be  obtained  ? 


CHAPTER  XIX 

THE  ALKALINE  EARTHS 

CALCIUM,  STRONTIUM,  BARIUM 

CALCIUM  (At.  Wt.  =40.1) 

THE  metals  which  we  have  thus  far  studied  are  all  univalent, 
or  their  ions  carry  one  electrical  charge  each.  In  the  second 
group  of  the  metals  the  ions  are  nearly  always  bivalent,  and 
in  the  calcium,  strontium,  barium  sub-group  they  are  always 
bivalent.  The  salts  which  these  elements  form  with  the  anions 
of  acids  are  of  the  general  type  MC12,  M(N03)2,  MS04,  and  so 
on.  With  this  conception  in  mind  we  may  now  proceed  to 
study  some  of  the  compounds  formed  by  the  several  members 
of  the  group. 

Occurrence,  Preparation,  and  Properties  of  Calcium.  —  Calcium 
occurs  very  widely  distributed  in  nature  and  in  large  quantities. 
The  carbonate  occurs  in  great  abundance  as  marble  if  well 
crystallized  and  pure,  or  if  impure  as  limestone  or  chalk.  If 
in  combination  with  magnesium  carbonate,  we  have  dolomite. 
Calcium  phosphate  occurs  in  considerable  quantity  in  certain 
phosphate  beds.  Gypsum,  or  calcium  sulphate,  occurs  in  con- 
siderable quantity,  while  calcium  fluoride,  or  fluor-spar,  exists 
in  certain  localities.  Calcium  comes  next  to  aluminium  and 
iron  in  the  order  of  abundance  in  the  earth. 

Calcium  is  best  prepared  by  decomposing  the  iodide  by 
metallic  sodium:  — 


Calcium  is  a  silvery-white  metal,  which  decomposes  water 
slowly  at  ordinary  temperatures. 

239 


240  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Calcium  Hydride,  CaH2,  is  formed  by  the  action  of  hydrogen 
on  the  hot  metal.  Strontium  and  barium  form  similar  com- 
pounds. 

Calcium  Oxide  or  Lime,  CaO.  —  Calcium  combines  with 
oxygen,  forming  the  oxide,  CaO.  The  oxide  is  most  conven- 
iently prepared  by  heating  the  carbonate:  — 

CaC03  =  C02  +  CaO. 

Calcium  oxide  is  a  white,  amorphous  powder,  which  is  ex- 
tensively used  in  a  number  of  chemical  operations. 

Calcium  oxide  does  not  melt  until  a  very  high  temperature 
is  reached  (about  3000°).  It  is,  therefore,  used  in  constructing 
the  Dmmmond  light.  When  the  flame  from  the  oxyhydrogen 
blowpipe  is  allowed  to  play  upon  a  cylinder  of  lime,  it  becomes 
highly  heated  and  at  this  high  temperature  gives  out  an 
enormous  amount  of  light. 

Calcium  Hydroxide  or  Slaked  Lime,  Ca(OH)2.  — When  lime  or 
calcium  oxide  is  thrown  into  water,  a  large  amount  of  heat  is 
evolved  and  calcium  hydroxide  is  formed  :  — 

CaO  +  H20  =  Ca(OH)2. 

This  process  is  known  as  slaking. 

Calcium  hydroxide  is  a  white  powder,  soluble  in  water  only 
to  the  extent  of  0.002  part  in  one  part  of  water.  This  solution 
is  known  as  lime-water.  A  mechanical  suspension  of  the  finely 
divided  calcium  hydroxide  in  water  is  known  as  milk  of  lime. 

Calcium  hydroxide  is  a  strongly  dissociated  compound.  It 
dissociates  thus:  — 

Ca(OH)2=Ca,0~H,OH. 

Its  solution  contains  aj  large  number  of  hydroxyl  ions,  and  it 
is,  therefore,  a  very  strong  base.  When  a  clear  solution  of 
calcium  hydroxide  is  allowed  to  stand  exposed  to  the  air  for  a 
short  time,  it  takes  up  carbon  dioxide  from  the  air,  forming 
flakes  of  the  insoluble  calcium  carbonate :  — 

Ca(OH)2  +  C02  =  CaC03  +  H20. 


THE  ALKALINE  EARTHS  241 

When  lime  is  exposed  to  the  air,  the  same  reaction  takes 
place  to  some  extent.  Calcium  oxide  takes  up  moisture  from 
the  air,  forming  the  hydroxide,  and  this  combines  in  part  with 
carbon  dioxide,  forming  the  carbonate.  The  white  powder 
formed  when  lime  is  exposed  to  the  air,  known  as  air-slaked 
lime,  is,  then,  a  mixture  of  calcium  oxide  and  calcium  carbon- 
ate. Lime  mixed  with  caustic  soda  is  known  as  soda-lime. 

Calcium  Chloride,  CaCl2,  6H20.  —  Calcium  chloride,  CaCl2, 
6  H20,  is  very  soluble  in  water,  producing,  when  dissolved,  a 
considerable  lowering  of  temperature.  By  mixing  this  salt  in 
the  proper  proportions  with  finely  powdered  ice  a  temperature 
as  low  as  —  30°  to  —  35°  can  be  produced.  The  salt  is  most 
readily  prepared  by  dissolving  marble  in  hydrochloric  acid, 
and  evaporating  the  solution  to  crystallization. 

On  account  of  its  attraction  for  water  anhydrous  calcium 
chloride  is  frequently  used  as  a  drying  agent,  especially  for 
gases.  These  are  passed  slowly  through  tubes  filled  loosely 
with  calcium  chloride,  and  most  of  the  water  is  removed  from 
the  gases  and  absorbed  by  the  chloride. 

Calcium  Hypochlorite,  Bleaching- powder,  Ca(OCl)2. — When 
chlorine  is  conducted  into  lime,  it  is  absorbed  by  the  lime.  The 
reaction  which  takes  place  may  be  represented  thus :  — 

2  Ca(OH)2  +  2  C12  -  CaCl2  +  Ca(OCl)2  +  2  H20,         (1) 

giving  a  mixture  of  calcium  chloride  and  hypochlorite ;  or  it 
may  be  represented  thus  :  — 

2  Ca(OH)2  +  2  C12  =  2  Ca  < ^+  2  H2O,  (2) 

forming  one  compound,  which  is  half  chloride  and  half  hypo- 
chlorite. 

It  is  difficult  to  decide  between  these  two  possibilities.  The 
fact  that  bleaching-powder  is  not  deliquescent,  while  calcium 
chloride  is  strongly  deliquescent,  would  indicate  that  there  is 
no  calcium  chloride  in  bleaching-powder.  When  treated  with 
an  acid  bleaching-powder  gives  up  all  of  its  chlorine :  — 


242  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Calcium  Sulphate,  CaS04.  —  The  sulphate  of  calcium  con- 
taining two  molecules  of  water  of  crystallization,  and  known  as 
gypsum  —  CaS04 .  2  H^O  —  occurs  abundantly  in  nature.  It 
dissolves  in  water  to  some  extent,  2  parts  in  1000,  and  is  fre- 
quently found  in  solution.  It  also  occurs  in  the  solid  form  in 
many  localities. 

Gypsum  is  useful  chiefly  on  account  of  the  transformations 
which  take  place  when  its  water  is  removed  by  heat,  and  the 
anhydrous  salt  is  brought  again  into  contact  with  water. 
When  gypsum  is  heated  to  107°  it  loses  a  part  of  its  water :  — 

2  CaS04 .  2  H20  =  2  CaS04 .  H20  +  3  H20. 

Gypsum  which  is  thus  partially  dehydrated  is  a  flowery 
powder,  and  is  known  as  plaster  of  par  is.  When  brought  in 
contact  with  water,  plaster  of  paris  takes  it  up  again  and  forms 
gypsum.  The  mass,  however,  is  now  finely  divided,  and 
hardens  after  a  few  minutes.  It  is  used  extensively  for 
making  mouldings  and  casts  of  objects,  especially  of  marble 
statuary. 

If  the  temperature  to  which  the  gypsum  is  heated  is  at  all 
high  (200°),  it  loses  all  of  its  water.  When  the  completely 
dehydrated  product  is  brought  in  contact  with  water,  it  com- 
bines with  the  water  very  slowly,  and  is  useless  as  far  as  mak- 
ing mouldings  is  concerned.  Such  gypsum  is  known  as  hard- 
burned  or  dead-burned  gypsum. 

Calcium  Carbide,  CaC2.  —  Calcium  carbide  has  come  into 
very  great  importance  recently  on  account  of  its  method  of 
preparation,  and  because  when  brought  into  the  presence  of 
water  it  readily  yields  the  illuminant  acetylene. 

The  carbide  of  calcium  is  prepared  by  heating  a  mixture 
of  finely  divided  carbon  and  lime  in  an  electric  furnace :  — 

3  C  +  CaO  =  CO  +  CaC2. 

Its  commercial  value  depends  entirely  upon  the  fact  that  it 
decomposes  with  water,  giving  acetylene  gas :  — 

CaC2  +  H20  =  CaO  +  C2H2. 


THE  ALKALINE  EARTHS  243 

Acetylene  gas  has  recently  come  into  great  prominence  as 
an  illuminant.  This  is  due  to  the  ease  with  which  it  can  be 
made  from  calcium  carbide.  Water  is  admitted  slowly  to  the 
carbide  when  the  acetylene  is  desired,  and  a  slow  or  rapid  cur- 
rent of  the  gas  generated  at  will. 

Calcium  Carbonate,  CaCQ3.  —  The  most  abundant  and  im- 
portant compound  of  calcium  which  occurs  in  nature  is  the 
carbonate.  Calcite  is  well-crystallized  calcium,  carbonate.  It 
occurs  in  beautifully  transparent,  crystalline  masses  in  Iceland, 
and  is  known  as  Iceland  spar.  Another  crystalline  variety  of 
calcium  carbonate  is  aragonite. 

Calcium  carbonate  occurs  in  great  abundance  in  less  beauti- 
fully crystallized  condition.  Marble  is  a  crystallized  form  of 
calcium  carbonate.  Limestone  is  also  calcium  carbonate,  and 
is  usually  crystalline,  but  the  crystals  are  generally  smaller 
than  in  marble  and  are  contaminated  with  various  impurities. 
Chalk  is  a  very  fine-grained  variety  of  calcium  carbonate, 
formed  chiefly  of  the  shells  of  microscopic  organisms.  In- 
deed, calcium  carbonate  is  frequently  of  organic  origin,  consist- 
ing of  shells  of  animals  which  have  been  more  or  less  meta- 
morphosed by  the  geological  processes  to  which  they  have  been 
subjected. 

Calcium  carbonate  can  be  readily  formed  in  the  laboratory 
by  treating  a  soluble  calcium  salt  with  a  soluble  carbonate :  — 

CaCl2  +  Na2C03  =  2  NaCl  +  CaC03. 

When  calcium  carbonate  is  heated,  it  undergoes  decomposi- 
tion into  lime  and  carbon  dioxide,  as  we  saw  when  we  were 
studying  lime :  — 

CaC03  =  CaO  +  C02. 

Primary  or  Acid  Calcium  Carbonate,  Ca(HC03)2,  is  formed 
when  the  normal  carbonate  is  dissolved  in  water  containing 
carbon  dioxide :  — 

CaC03  +  H20  +  C02  =  Ca(HC03)2. 


244  ELEMENTS  OF   INORGANIC   CHEMISTRY 

Into  water  containing  normal  calcium  carbonate  in  suspen- 
sion conduct  a  current  of  carbon  dioxide,  and  the  calcium 
carbonate  will  pass  into  solution  as  the  acid  carbonate.  The 
acid  carbonate  cannot,  however,  be  isolated.  Indeed,  when  its 
solution  is  boiled,  the  acid  carbonate  is  decomposed  into  the 
normal  carbonate  arid  carbon  dioxide  is  given  off : 

Ca(HC03)2  =  CaC03  +  H20  +  C02. 

When  a  solution  of  the  acid  carbonate  is  evaporated  at  ordi- 
nary temperatures,  the  normal  carbonate  is  deposited.  This  is 
the  way  in  which  the  stalagma  in  caverns  are  formed.  The 
waters  containing  acid  calcium  carbonate  in  solution  percolate 
through  the  roof  of  a  cavern,  and,  evaporating,  deposit  calcium 
carbonate  on  the  ceiling.  This  continues,  the  solution  of  the 
acid  carbonate  trickling  over  the  outside  of  the  deposit  already 
formed,  until  frequently  very  beautiful  hanging  columns  result. 
Such  formations  suspended  from  the  ceiling  or  sides  of  a 
cavern  are  called  stalactites.  The  solution  of  the  acid  carbonate 
frequently  drops  off  of  the  stalactite,  since  the  rate  of  evapora- 
tion in  a  closed  space  beneath  the  surface  of  the  earth  must  be 
very  slow.  Where  it  drops  on  the  floor  of  the  cavern  it  evapo- 
rates and  deposits  the  carbonate,  and  we  frequently  find  col- 
umns and  pillars  of  calcium  carbonate  growing  upwards  from 
the  floor  of  caverns.  Such  growths  are  known  as  stalagmites. 

Natural  waters  frequently  contain  carbon  dioxide  in  consid- 
erable quantity.  This  is  produced  by  decomposing  vegetable 
matter  in  the  soil  through  which  they  percolate,  and  is  also 
taken  from  the  atmospheric  air.  Consequently,  many  natural 
waters  contain  calcium  carbonate  in  solution.  Such  are  known 
as  hard  waters. 

Phosphates  of  Calcium.  —  The  three  calcium  salts  of  phos- 
phoric acid  are  all  known.  They  are  the  tricalcium  phosphate 
or  normal  salt,  Ca3(P04)2,  the  secondary  salt,  CaHP04,  and  the 
primary  salt,  Ca(H.2P04)2. 

Tricalcium  phosphate,  Ca3(P04)2,  is  found  in  large  quantity 
in  the  bones  of  animals,  and  is  therefore  very  important  in 


THE   ALKALINE    EARTHS  245 

connection  with  animal  life.  When  bones  are  heated  to  a  high 
temperature  in  contact  with  the.  air,  the  organic  matter  is 
destroyed,  and  the  calcium  phosphate  and  other  mineral  matter 
in  the  bones  remain  behind  in  the  bone-ash. 

The  normal  calcium  phosphate  also  occurs  in  nature  as  phos- 
phorite, and  in  combination  with  calcium  chloride  or  fluoride  as 
apatite. 

The    phosphoric    ions  —  tertiary,    secondary,    and   primary, 

P04,  HP04,  or  H2F04  —  are  of  fundamental  importance  for  the 
growth,  and  especially  for  the  seeding,  of  plants.  Among 
these  are  the  very  valuable  cereals  wheat  and  corn.  These 
plants  gradually  remove  the  phosphoric  acid  ion  from  the  soil, 
and  the  latter  would  soon  become  impoverished  in  this  sub- 
stance were  it  not  supplied  to  the  soil  artificially.  The  most 
important  constituent  of  commercial  fertilizer  is  phosphoric 
acid  ions. 

Secondary  calcium  phosphate,  CaHP04,  is  formed  by  the 
action  of  a  soluble  calcium  salt  on  disodium  phosphate  in  the 
presence  of  a  little  acid. 

If  a  little  acetic  acid  is  added  to  the  solution  of  disodium 
phosphate  and  calcium  chloride  then  introduced,  the  tricalcium 
phosphate  is  formed  at  once  :  — 

3  CaCl2  +  2  NagHPO,  +  (CH3COOH) 
=  Ca3(P04)2  +  4  NaCl  +  2  HC1  +  (CH3COOH). 

If  tricalcium  phosphate  is  treated  with  a  strong  acid,  i.e. 
with  a  concentrated  solution  of  hydrogen  ions  in  the  proper 
proportions,  the  primary  calcium  phosphate  is  formed  :  — 

Ca3(P04)2  +  2  H2S04  =  2  CaS04  +  Ca(H2P04)2. 

This  is  the  commercial  superphosphate. 

In  preparing  commercial  fertilizer  the  tertiary  calcium  phos- 
phate is  the  starting-point  in  obtaining  phosphoric  acid  ions. 

The  tricalcium  phosphate  in  ground  bone,  or  in  finely  ground 
phosphate  rock,  is  treated  with  sulphuric  acid  and  converted 
into  phosphates  which  are  somewhat  soluble  in  water.  This 


246  ELEMENTS   OF  INORGANIC   CHEMISTRY 

is  known  as  soluble  or  available  phosphoric  acid,  while  the 
phosphoric  acid  in  the  form  of  tricalciuni  phosphate  is  known 
as  insoluble  phosphoric  acid. 

Calcium  Silicate,  CaSiOo. — The  silicate  of  calcium  occurs  in 

o 

nature  in  such  well-known  minerals  as  mica  and  garnet.  Its 
chief  importance  is  in  connection  with  the  manufacture  of  glass. 
Glass  is,  in  general,  an  amorphous  mixture  of  the  silicate  of 
calcium  with  the  silicates  of  the  alkalies.  There  are  a  number 
of  varieties  of  glass,  and  a  few  of  these  will  be  reconsidered. 

Glass  is  made  by  fusing  together  sand,  calcium  carbonate, 
and  the  carbonate  of  the  alkali  desired.  If  sodium  carbonate 
is  used,  we  have  soda  glass;  if  potassium  carbonate  is  employed, 
potash  glass,  etc. 

Soda  glass  is  prepared  by  fusing  together  sodium  carbonate, 
calcium  carbonate,  and  silicon  dioxide.  Soda  glass  is  readily 
fusible,  and  is  easily  attacked  by  chemical  reagents  such  as 
boiling  alkalies.  It  is  blown  into  cylinders,  which  are  then 
opened  and  flattened,  and  cut  into  ordinary  window  panes. 
This  is  known  as  soft  glass,  because  it  is  easily  worked  in 
the  blast-lamp,  and  has  applications  in  the  chemical  laboratory, 
although,  on  account  of  its  solubility,  it  is  not  well  adapted  for 
bottles  for  holding  chemical  reagents. 

Bohemian  or  potassium  glass  is  a  potassium  calcium  silicate. 
It  is  much  harder  than  the  soda  glass,  fuses  at  a  much  higher 
temperature,  and  is  much  more  resistant  to  chemical  reagents. 
It  is,  therefore,  valuable  to  the  chemist,  and  is  extensively 
employed  in  the  manufacture  of  apparatus  which  is  to  be  heated 
to  a  high  temperature,  such  as  combustion-tubing  and  the  like. 

Flint-glass  consists  of  potassium  and  lead  silicates,  the  lead 
taking  the  place  of  calcium  in  ordinary  soda  or  potassium  glass. 
It  has  such  a  high  refractive  power  that  it  is  used  in  making 
optical  lenses. 

The  colored  glasses  are  prepared  by  adding  to  the  fused  sili- 
cates, oxides  of  certain  metals  which  give  the  desired  color  to 
the  glass.  Yellow  glasses  owe  their  color  to  uranium  or  anti- 
mony; blue  glasses  to  cobalt  or  manganese ;  red  glasses  to  copper, 


THE   ALKALINE   EARTHS  247 

iron,  or  sometimes  gold ;  green  glasses  to  chromium  or  copper, 
and  so  on.  Glasses  of  almost  every  shade  of  color  have  been 
prepared  by  using  different  coloring  constituents  or  mixtures  of 
these  constituents. 

Most  of  the  glass  objects  with  which  we  are  ordinarily  fa- 
miliar are  blown.  A.  large  ball  of  molten  glass  is  taken  on  the 
end  of  a  hollow  metal  tube,  through  which  the  breath  can  be 
blown.  The  tube  is  then  moved  rapidly  backwards  and  for- 
wards through  the  air  beneath  the  glass-blower,  who  drives  air 
through  the  tube  at  the  desired  rate  and  time.  The  glass  usually 
takes  the  form  of  a  hollow  cylinder,  which  is  blown  out  to 
the  desired  thickness.  This  is  cracked,  flattened,  and  cut  into 
the  desired  size.  In  this  way  flat  panes  of  glass  are  made. 
Bottles  are  blown  into  moulds,  and  other  glass  objects  of  definite 
shape  are  either  blown  into  moulds  or  moulded.  Such  objects 
must  be  annealed. 

The  properties  of  the  glass  is  largely  conditioned  by  the  way 
in  which  it  is  annealed.  If  glass  is  cooled  with  moderate 
rapidity,  it  has  the  properties  which  we  ordinarily  associate  with 
it.  If  cooled  very  rapidly,  however,  it  has  very  different  prop- 
erties. It  is  under  a  considerable  strain,  and  when  the  surface 
is  fractured  in  any  way,  the  glass  flies  in  pieces  almost  with 
explosive  violence. 

The  Prince  Rupert  drops,  made  by  dropping  molten  glass 
into  water,  are  examples  of  this  condition. 

On  the  other  hand,  when  glass  is  cooled  very  slowly,  as  by 
introducing  it  when  hot  into  hot  oil,  or  by  placing  it  in  an  oven 
which  is  cooled  slowly,  it  is  much  less  easily  broken  than 
ordinary  glass.  It  acquires  considerable  elasticity,  and  can  be 
struck  a  fairly  hard  blow  without  injury. 

STRONTIUM  (At.  Wt.  =  87.68) 

The  element  strontium  resembles  calcium  very  closely  in  its 
properties,  and  in  the  properties  of  its  compounds.  It  will, 
therefore,  be  treated  very  briefly,  certain  differences  between 
the  two  being  pointed  out. 


248  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Occurrence,   Preparation,  and    Properties  of   Strontium.  — 

Stontium  occurs  in  nature  chiefly  in  the  form  of  two  salts 
which  are  well-known  minerals.  These  are  strontium  carbon- 
ate or  strontianite,  and  strontium  sulphate  or  celestite. 

The  element  is  prepared  most  conveniently  by  electrolyzing 
the  fused  chloride. 

Strontium  resembles  calcium  in  its  appearance  and  properties. 

++ 
Salts  of  Strontium.  —  Strontium  forms  the  bivalent  ion  Sr, 

++ 
which  is  strictly  analogous  to  the  calcium  ion  Ca.     It  combines 

with  two  hydroxyl  ions,  forming  strontium  hydroxide,  Sr(OH)2, 
whose  aqueous  solution  is  strongly  basic.  This  substance  is 
more  soluble  in  water  than  calcium  hydroxide,  and  crystal- 
lizes from  the  aqueous  solution  with  eight  molecules  of  water : 
Sr(OH)2-8  H20.  Strontium  nitrate,  Sr(N03)2,  like  strontium 
salts  in  general,  gives  a  beautiful  red  color  to  a  colorless  flame. 
It  is  used,  because  of  this  property,  to  produce  red  light  in  fire- 
works and  other  pyrotechnic  displays. 

The  strontium  ion  combines  with  the  anions  of  acids,  forming 
in  general  the  same  insoluble  compounds  as  calcium.  Certain 
differences  in  the  degree  of  solubility,  however,  manifest  them- 
selves. Strontium  combines  with  the  ion  S04,  forming  stron- 
tium sulphate,  SrS04.  This  salt  occurs  in  nature  as  celestite, 
and  since  it  is  only  slightly  soluble  in  water,  is  formed  when  a 
soluble  sulphate  is  added  to  a  soluble  strontium  salt :  — 

Na2S04  +  SrCl,  =  2  NaCl  +  SrS04. 

Strontium  sulphate  is  much  less  soluble  in  water  than  cal- 
cium sulphate,  and  is  practically  insoluble  in  a  mixture  of 
water  and  alcohol. 

BARIUM  (At.  Wt.=  137.4) 

An  element  closely  allied  to  calcium  and  strontium  is  ba- 
rium. This  element  occurs  chiefly  as  the  sulphate,  BaS04, 
which  is  known  as  barite  or  heavy  spar  ;  and  as  the  carbonate, 
BaC03,  known  as  ivitherite. 


THE   ALKALINE    EARTHS  249 

Barium,  like  strontium  and  calcium,  is  prepared  by  electro- 
lyzing  the  fused  chloride.  The  metal  barium  is  white,  takes 
up  oxygen  from  the  air,  and  decomposes  water  with  evolution 
of  hydrogen. 

Oxides  of  Barium.  —  Barium  forms  two  oxides  —  the  normal 
oxide,  BaO,  and  the  dioxide,  Ba02.  Barium  oxide,  BaO,  is 
formed  by  heating  the  nitrate:  — 

2  Ba(N03)2  =  2  BaO  +  4  N02  +  Oa. 

Barium  dioxide,  BaG2,  is  formed  by  heating  the  oxide  to  500° 
in  a  current  of  air  or  oxygen  :  — 

2  BaO  +  02  =  2  Ba02. 

Barium  dioxide  is  an  excellent  "  carrier  of  oxygen,"  since  at 
a  somewhat  higher  temperature  it  gives  up  its  excess  of  oxygen 
and  forms  barium  oxide  again  :  — 

2Ba02  =  2BaO  +  02. 

We  have  already  become  familiar  with  this  substance  in 
connection  with  the  preparation  of  hydrogen  dioxide.  When 
it  is  treated  with  an  acid,  the  following  reaction  takes  place  :  — 

BaO2  +  2  HC1  =  BaCl2  +  H202. 

Barium  Hydroxide,  Ba(OH)2.  —  The  hydroxide  of  barium  is 
formed  when  the  oxide  is  dissolved  in  water  :  — 


20  =  Ba(OH)2. 

Barium  oxide  is  much  more  soluble  in  water  than  strontium 
oxide,  which  in  turn  is  more  soluble  than  calcium  oxide. 

Barium  Sulphate,  BaS04.  —  The  sulphate  of  barium,  or  heavy 
spar,  occurs  in  nature  as  stated  above.  It  is  really  formed 
whenever  a  soluble  sulphate  is  added  to  a  soluble  barium  salt  :  — 

Ba(N03)2  +  K2S04  =  2  KN03  -f  BaS04. 

It  is  the  most  insoluble  sulphate  known,  and  is,  therefore, 
the  form  in  which  sulphuric  acid  is  precipitated  and  weighed 
in  quantitative  determinations  of  this  acid. 


250  ELEMENTS   OF   INORGANIC   CHEMISTRY 

One  decomposition  of  barium  sulphate  is  of  more  than 
ordinary  interest.  When  the  sulphate  is  boiled  with  a  solution 
of  sodium  carbonate,  it  is  transformed  in  part  into  barium 
carbonate  :  — 

BaS04  +  Na2C03  =Na2S04  +  BaC03. 

This  transformation  is  at  first  only  partial.  If  the  solution 
of  sodium  carbonate  and  sodium  sulphate  is  poured  off  after 
a  time,  and  a  new  solution  of  sodium  carbonate  added,  the 
decomposition  of  the  barium  sulphate  into  carbonate  will  pro- 
ceed farther.  By  repeating  this  for  a  few  times,  practically  all 
of  the  barium  sulphate  can  be  transformed  into  carbonate. 
This  is  one  of  the  very  best  examples  of  the  effect  of  mass  on 
chemical  activity.  By  renewing  the  solution  of  sodium  carbonate, 
i.e.  by  increasing  its  mass,  and  by  pouring  off  the  solution  of 
sodium  sulphate  formed,  i.e.  by  decreasing  its  mass,  the  trans- 
formation in  the  sense  of  the  above  equation  can  be  made 
practically  complete. 

The  Law  of  Mass  Action.  —  We  have  seen  a  number  of 
examples  of  the  effect  of  mass  on  chemical  activity.  The 
effect  has  been  studied  exhaustively  in  a  great  many  cases,  and 
the  action  of  mass  so  well  worked  out  that  it  can  be  formulated 
in  a  simple  algebraic  expression.  If  we  represent  the  quantities 
of  two  substances  in  solution  which  are  reacting  chemically  by 
p  and  q,  and  the  quantities  of  the  two  substances  formed  by  pl 
and  qlt  the  following  relation  obtains  :  — 


in  which  k  and  \  are  two  constants  whose  values  depend  upon 
the  nature  of  the  substances  in  question. 

This  generalization,  whose  significance  cannot  at  this  stage 
in  the  study  of  chemistry  be  fully  appreciated,  is  of  funda- 
mental importance  in  dealing  with  the  nature  of  chemical  re- 
actions and  with  the  velocities  with  which  they  take  place. 

Relations  between  Calcium,  Strontium,  and  Barium.  —  The 
three  alkaline  earth  metals  resemble  one  another  closely  in 


THE   ALKALINE   EARTHS  .  251 

their  chemical  properties.  Certain  differences,  however,  mani- 
fest themselves.  The  different  solubilities  of  the  salts  of  these 
elements  with  a  given  acid  are  very  important,  especially  in 
connection  with  the  separation  of  these  elements  from  one 
another. 

Take  the  hydroxides  of  the  three  elements :  Calcium  hydrox- 
ide is  the  least  soluble ;  strontium  hydroxide  is  more  soluble 
than  calcium;  while  barium  hydroxide  is  still  more  soluble 
than  strontium  hydroxide. 

When  we  turn  to  the  sulphates,  we  find  exactly  the  opposite 
relations.  Barium  sulphate  is  the  most  insoluble  of  the  three ; 
then  comes  strontium  sulphate,  and  finally  calcium  sulphate. 

All  three  of  these  elements  form  insoluble  carbonates,  while 
barium  alone  forms  an  insoluble  chromate. 

These  differences  in  solubility  are  made  use  of  to  detect  the 
alkaline  earths  when  in  solution  in  the  presence  of  one  another. 

The  Spectroscopic  Study  of  the  Alkalies  and  Alkaline  Earths. 
—  When  white  light  is  passed  through  a  glass  prism,  it  is  bent 
out  of  its  course,  or  refracted,  as  we  say.  White  light  is  not 
homogeneous,  but  is  composed  of  several  different  colors. 
Light  of  one  color  is  bent  to  a  different  extent  from  light  of 
another  color.  This  can  be  seen  by  allowing  white  light  to  fall 
on  a  prism  of  glass,  and  placing  the  eye  on  the  opposite  side 
of  the  prism.  We  would  see  a  number  of  colors  arranged  in 
a  definite  order  from  red  through  yellow  to  green  and  blue. 
These  are  known  as  the  prismatic  colors. 

An  instrument  has  been  devised  for  breaking  up  light  into 
its  several  colors.  This  is  known  as  the  spectroscope.  It  con- 
sists essentially  of  a  prism,  a  tube  containing  at  one  end  a  nar- 
row slit  through  which  the  incident  light  enters,  and  another 
tube  through  which  the  refracted  light  is  seen. 

If  light  of  a  definite  color  is  placed  in  front  of  the  spectro- 
scope, it  will  appear  in  the  instrument  as  a  bright  line  or  bright 
lines.  When  sodium  light  is  viewed  through  a  spectroscope, 
a  bright  yellow  line  will  appear. 

The  spectroscope  should  also  be  provided  with  a  scale,  so 


252  ELEMENTS   OF   INORGANIC   CHEMISTRY 

that  the  exact  position  of  any  line  can  be  determined.  The 
line  produced  by  an  element  has  a  perfectly  definite  and  fixed 
position,  and  can  be  used  to  determine  whether  the  element  is 
present.  This  principle  is  utilized  in  spectroscopic  analysis. 
Where  several  elements  are  present,  all  of  which  give  colored 
flames,  it  is  impossible  by  examining  the  flame  with  the  eye 
alone  to  tell  which  is  present.  Knowing  the  exact  position  in 
the  spectroscope  of  the  line  (or  lines)  produced  by  the  different 
elements,  it  is  only  necessary  to  throw  the  light  from  a  flame 
containing  any  number  of  elements  upon  a  spectroscope,  in 
order  to  read  off  at  once  what  elements  are  present. 

The  spectroscope  has  been  of  the  greatest  use  to  the  chemist 
in  identifying  substances,  in  determining  whether  they  were 
pure,  and  in  the  discovery  of  new  elements. 

EXPERIMENTS  WITH   CALCIUM,   STRONTIUM,  AND 
BARIUM 

Experiment  150.  Lime  formed  by  Heating  Calcium  Car- 
bonate  (A  small  piece  of  marble;  a  piece  of  charcoal; 

blowpipe.) 

Heat  a  piece  of  marble  upon  a  piece  of  charcoaljvvith  the  tip 
of  the  flame  of  the  blowpipe.  After  heating  the  marble  for 
several  minutes  throw  it  into  water.  Test  the  reaction  of  the 
water  with  red  litmus  paper.  (See  page  240.) 

Experiment  151.  Slaking  of  Lime;  Properties  of  Cal- 
cium Hydroxide.  —  (Freshly  burned  lime  ;  soluble  salts  of  the 
same  heavy  metals  as  in  Experiment  143.) 

Treat  50  grams  of  freshly  burned  lime  in  a  porcelain  dish 
with  a  little  wrater.  The  water  is  taken  up  and  a  large  amount 
of  heat  is  evolved.  The  lime  crumbles  and  falls  to  pieces, 
yielding  a  fine  powder.  This  is  slaked  lime  or  calcium  hydrox- 
ide. Treat  a  part  of  this  powder  with  a  litre  or  two  of  water. 
Allow  the  solution  to  settle  and  decant  the  clear  liquid.  This 
is  lime-water,  or  calcium  hydroxide,  in  solution. 

Calcium  hydroxide  has  properties  closely  allied  to  those  of 
potassium  hydroxide.  Treat  solutions  of  salts  of  the  heavy 
metals  with  lime-water ;  the  hydroxides  of  the  heavy  metals 
are  thrown  down.  Compare  the  results  with  those  obtained  in 
Experiment  143.  (See  page  240.) 


THE  ALKALINE   EARTHS  253 

Experiment  152.  Preparation  and  Properties  of  Calcium 
Chloride.  —  (Marble ;  hydrochloric  acid ;  thermometer.) 

Dissolve  marble  in  hydrochloric  acid,  using  an  excess  of 
marble.  Evaporate  the  solution  and  calcium  chloride  will 
crystallize  out.  Mix  the  powdered  crystals  of  calcium  chloride 
with  one-third  their  volume  of  sodium  chloride.  Plunge 
a  thermometer  into  the  mixture.  What  evidence  have  you 
that  there  is  a  marked  lowering  of  temperature  ?  (See  page 
240.) 

Experiment  153.     Gypsum ;  Plaster  of  Paris.  —  (Gypsum.) 

Heat  a  little  gypsum  to  140°  for  a  number  of  hours.  It  is 
partly  dehydrated,  and  forms  a  flowery  powder  —  plaster  of 
paris.  Add  a  little  water  to  plaster  of  paris.  It  forms  a 
pasty  mass,  which  rapidly  hardens  and  can  be  used  for  mould- 
ing objects. 

Heat  a  little  gypsum  to  200°  for  an  hour  or  two.  Add  water. 
Does  it  behave  like  plaster  of  paris  ?  What  is  the  difference  ? 
(See  page  242.) 

Experiment  154.      Calcium  Carbide.  —  (Calcium  carbide.) 

Introduce  a  small  piece  of  calcium  carbide  into  a  test-tube 
and  drop  water  slowly  upon  it.  Smell  the  odor  of  the  escap- 
ing gas.  Does  it  burn  ?  What  remains  in  the  test-tube  ? 
Add  a  little  of  the  solution  to  solutions  of  salts  of  the  heavy 
metals.  What  are  precipitated?  (See  page  242.) 

Experiment  155.  Calcium  Carbonate  and  Acid  Calcium 
Carbonate (Lime-water;  carbon  dioxide  generator.) 

Pass  carbon  dioxide  into  lime-water.  Insoluble  calcium 
carbonate  will  be  precipitated.  WTrite  the  equation  ? 

Continue  to  pass  carbon  dioxide  into  water  containing  cal- 
cium carbonate  in  suspension.  The  calcium  carbonate  will 
pass  into  solution  again.  Write  equation  and  explain  the  fact  ? 

Boil  the  clear  solution ;  calcium  carbonate  will  be  precipi- 
tated again.  Explain  ?  Write  equation  ? 

Pass  carbon  dioxide  again  through  the  solution  containing 
calcium  carbonate  in  suspension.  It  will  redissolve  ?  (See 
page  243.) 

Experiment  156.  Flame  Reaction  of  Calcium —  (Calcium 
compound;  platinum  wire.) 

Dip  a  clean  platinum  wire  into  a  calcium  compound  and 
insert  the  wire  into  the  colorless  flame  of  a  bunsen  burner. 
Note  the  orange-red  color  imparted  to  the  flame  by  the  calcium 
compound.  This  flame  is  characteristic  of  calcium. 


254  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Experiment  157.  Flame  Reaction  of  Strontium.  —  (Stron- 
tium compound;  platinum  wire.) 

Introduce  a  strontium  compound  on  a  clean  platinum  wire 
into  a  bunsen  flame.  Note  the  deep  red  color  imparted  to  the 
flame.  This  flame  reaction  is  characteristic  of  strontium. 

Experiment  158.  Properties  of  Barium  Hydroxide.  — 
(Baryta  water  ;  soluble  carbonates  ;  soluble  salts  of  the  heavy 
metals.) 

Treat  a  number  of  soluble  carbonates  with  baryta  water. 
White  barium  carbonate  is  precipitated.  Add  acid  to  barium 
carbonate.  What  takes  place  ? 

Treat  solutions  of  salts  of  the  heavy  metals  (lead  nitrate; 
zinc  chloride ;  cadmium  chloride  ;  aluminium  chloride ;  chro- 
mium chloride,  etc.)  with  baryta  water.  What  is  formed  ? 

Barium  hydroxide,  like  strontium  hydroxide,  behaves  like 
calcium  hydroxide.  They  all  resemble  chemically  the  hydrox- 
ides of  the  alkalies.  (See  page  249.) 

Experiment  159.  Barium  Sulphate  decomposed  by  Sodium 
Carbonate.  —  (Barium  chloride ;  sodium  sulphate ;  sodium 
carbonate.) 

To  a  solution  of  barium  chloride  add  a  solution  of  sodium 
sulphate.  Insoluble  barium  sulphate  is  precipitated.  Filter 
this  off  and  wash  thoroughly  with  hot  water. 

Transfer  the  barium  sulphate  to  a  small  beaker  and  pour  over 
it  a  solution  of  sodium  carbonate.  Boil  for  ten  or  fifteen 
minutes.  Then  filter  off  the  solution  and  test  for  sulphuric 
acid  in  the  filtrate,  by  adding  barium  chloride  and  hydrochloric 
acid.  Hydrochloric  acid  is  added  to  prevent  barium  carbonate 
from  being  precipitated.  Explain  ? 

Sulphuric  anions  will  be  found  in  the  filtrate  due  to  the  for- 
mation of  sodium  sulphate,  which  is  soluble.  (See  page  250.) 

Experiment  160.  Flame  Reaction  of  Barium.  —  (Barium 
compound;  platinum  wire.) 

Introduce  a  barium  compound  on  a  clean  platinum  wire  into  a 
colorless  bunsen  flame.  A  bright  green  color  will  be  imparted  to 
the  flame.  This  is  the  characteristic  flame  reaction  of  barium. 

Experiment  161.  Spectroscopic  Examination  of  the  Alkalies 
and  Alkaline  Earths.  —  (Sodium  salt ;  potassium  salt ;  calcium 
salt ;  strontium  salt ;  barium  salt ;  fine  pieces  of  platinum 
wire ;  spectroscope.) 

Clean  the  platinum  wires  by  dipping  them  into  hydrochloric 
acid,  and  plunge  one  into  each  of  the  above  salts. 


THE   ALKALINE   EARTHS  255 

Place  these  wires,  one  at  a  time,  in  the  colorless  flame  in 
front  of  the  spectroscope.  Observe  the  bright  lines  which 
flash  out  in  each  case.  Carefully  note  the  color  and  position 
of  the  lines  in  the  spectroscope  in  each  case.  The  exact  posi- 
tion of  the  lines  is  characteristic  of  the  element  in  question. 
(See  page  251.) 


PROBLEMS 

1.  What  weight   of   lime   and  of  carbon   dioxide   can  be 
obtained  from  1  kilogram  of  marble  ?     How  much  bleaching- 
powder  from  the  above  amount  of  lime  ? 

2.  Weight  of  acetylene  obtainable  from  the  calcium  carbide 
which  can  be  formed  from  1  gram  of  carbon  ? 

3.  Amount  of  strontium  nitrate  which  could  be  formed  from 
100  grams  of  strontium  hydroxide  ?     Percentage  of  strontium 
in  each  substance  ? 

4.  Weight   of  the  barium  sulphate  which  would  be  pre- 
cipitated when  sulphuric  acid  is  added  to  a  solution  contain- 
ing 0.75  gram  of  barium  nitrate  ? 


CHAPTER  XX 

THE  MAGNESIUM   GROUP 
GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY 

GLUCINUM  (At.  Wt.  =  9.1) 

THE  first  two  members  of  this  group,  glucinum  and  mag- 
nesium, correspond  rather  closely  in  their  properties  to  the 
elements  of  the  calcium  group.  The  first  of  these  elements, 
glucinum  or  beryllium,  is  comparatively  rare.  It  occurs  chiefly 
in  the  mineral  beryl,  which  is  a  silicate  of  aluminium  and 
beryllium. 

MAGNESIUM  (At.  Wt.  =  24.36) 

The  second  member  of  the  series  in  the  order  of  increasing 
atomic  weights,  but  the  first  member  in  the  order  of  importance, 
and  the  one  from  which  the  series  takes  its  name,  is  magnesium. 
Magnesium  occurs  in  nature  in  the  form  of  several  salts.  Mag- 
nesium carbonate,  magnesite,  MgC03,  is  fairly  widely  distrib- 
uted, while  the  double  carbonate  of  magnesium  and  calcium, 
dolomite,  forms  whole  mountain  ranges.  Magnesium  also  occurs 
as  the  sulphate,  MgS04.  H20,  —  kieserite. 

Carnallite,  —  the  double  chloride  of  potassium  and  magnesium 

—  KMgCl3.6  H20,  is  also  found  in  large  quantities  in  certain 

salt  deposits.     The  double  silicates  of  magnesium  and  other 

metals  constitute  some  of  the  best-known  minerals,  such  as 

hornblende,  serpentine,  talc,  asbestos,  etc. 

Magnesium  is  prepared  by  electrolyzing  fused,  anhydrous 
carnallite.  When  heated  it  takes  fire  and  burns  with  a  brill- 
iant white  light,  which,  on  account  of  its  richness  in  light  of 
short  wave-lengths,  is  frequently  used  for  illuminating  purposes 
in  photography. 

256 


THE   MAGNESIUM   GROUP  257 

Magnesium  Oxide,  MgO,  and  Magnesium  Hydroxide,  Mg(OH)2. 

—  Magnesium  oxide,  or  magnesia,  is  formed  when  magnesium 
is  burned  in  the  air,  or  when  the  carbonate  or  hydroxide  of 
magnesium  is  heated.     Magnesia  melts  only  at  enormously 
high  temperatures.     It  is,  therefore,  used  to  line  vessels  which 
are  subjected  to  high  temperatures. 

Magnesium  oxide  dissolves  in  water  to  only  a  slight  extent, 
"forming,  however,  magnesium  hydroxide. 

Magnesium  Chloride,  MgCl2 .  6  H20.  —  Magnesium  chloride, 
like  the  chlorides  of  calcium  and  strontium,  contains  six 
molecules  of  water  of  crystallization.  It  occurs  in  nature  in 
combination  with  potassium  chloride  as  carnallite,  and  can 
readily  be  prepared  by  dissolving  magnesium  carbonate  or 
oxide  in  hydrochloric  acid. 

Magnesium  Sulphate,  MgS04 .  7  H20. — Magnesium  sulphate 
containing  one  molecule  of  water,  occurs  in  nature  as  the 
mineral  kieserite,  MgS04 .  H20.  The  ordinary  salt,  known  as 
Epsom  salt  since  it  is  contained  in  the  Epsom  springs,  contains 
seven  molecules  of  water,  MgS04 .  7  H20. 

Magnesium  sulphate  is  used  in  medicine  as  a  purgative. 

Phosphates  of  Magnesium.  —  Magnesium  forms  the  primary 
Mg(H2P04)2;  the  secondary,  MgHP04;  and  the  tertiary,  or 
normal,  phosphate,  Mg3(P04)2.  These  resemble  the  phosphates 
of  calcium  so  closely  that  a  detailed  discussion  is  unnecessary. 
One  phosphate  of  magnesium,  however,  is  of  special  interest. 
This  is  the  ammonium  magnesium  phosphate  already  mentioned 

—  NH4MgP04.6H20. 

When  heated  it  decomposes  as  follows :  — 

2  NH4MgP04  =  2  NH3  +  H20  +  Mg2P207, 

yielding  magnesium  pyrophosphate,  which  is  a  very  stable  sub- 
stance and  of  constant  composition.  It  is,  therefore,  an  excellent 
form  in  which  to  weigh  either  phosphoric  acid  or  magnesium. 

Magnesium  Nitride,  Mg3N2,  is  formed  by  the  direct  action  of 
nitrogen  on  red-hot  magnesium.  Calcium  and  barium  also 
form  nitrides  quite  readily. 


258  ELEMENTS  OF  INORGANIC   CHEMISTRY 

ZINC  (At.  Wt.  =  65.4) 

An  element  closely  allied  to  magnesium  in  many  of  its  prop- 
erties, but  differing  from  it  markedly  in  others,  is  zinc.  This 
well-known  element  occurs  in  nature  in  abundance  in  a  number 
of  compounds.  Zinc  blende,  which  is  the  sulphide,  ZnS,  is  one 
of  the  most  common.  It  also  occurs  as  the  carbonate,  ZnC03, 
smitksonite,  and  the  silicate,  Zn2Si04,  calcimine. 

Zinc  is  prepared  by  reducing  with  carbon  the  oxide,  which  is 
usually  obtained  by  roasting  the  sulphide  in  the  air :  — 
ZnO  +  C  =  CO  +  Zn. 

The  metal  boils  at  950°,  and  is,  therefore,  easily  distilled  off 
when  once  set  free.  It  condenses  in  the  form  of  a  fine  powder 
or  dust,  which  is  known  as  zinc  dust.  The  metal  can  be  melted 
and  cast  into  sticks,  or  poured  into  water  while  molten  and 
produces  granulated  zinc. 

Zinc  is  used  for  covering  objects  which  if  exposed  to  air  or 
water  would  rust.  Iron  objects  thus  protected  are  known  as 
galvanized.  Iron  objects  are  galvanized  by  dipping  them  into 
molten  zinc.  Zinc  is  also  extensively  used  in  constructing 
primary  cells.  The  chief  uses  of  zinc,  however,  are  as  alloys 
with  other  metals.  The  best  known  of  these  is  brass,  an  alloy 
with  copper.  When  nickel  is  added,  we  have  the  alloy  of  zinc, 
copper,  and  nickel,  known  as  German  silver. 

Zinc  readily  combines  with  mercury  when  the  clean  surface 
of  the  metal  is  brought  in  contact  with  the  mercury,  forming 
zinc  amalgam. 

Zinc  Oxide,  ZnO,  and  Hydroxide,  Zn(OH)2.  —  Zinc  oxide  is 
formed  when  zinc  burns  in  the  air.  It  is  known  from  its 
general  appearance  as  philosopher's  wool.  It  is  prepared  more 
conveniently  by  heating  basic  zinc  carbonate.  This  loses  car- 
bon dioxide  and  water,  and  zinc  oxide  remains  behind.  Zinc 
oxide  is  used  as  a  pigment  under  the  name  of  zinc  white. 

Zinc  hydroxide,  Zn(OH)2,  is  formed  when  a  soluble  zinc  salt 
is  treated  with  a  small  amount  of  an  alkaline  hydroxide :  — 
ZnCl2  -f  2  KOH  =  2  KC1  +  Zn(OH)2. 


THE  MAGNESIUM  GROUP  259 

Zinc  hydroxide  is  a  white,  flocculent  precipitate,  which  dis- 
solves readily  in  acids  forming  the  corresponding  salts.  It  is 
therefore  basic  with  respect  to  acids,  and  as  it  dissolves,  must 
dissociate  into  zinc  and  hydroxyl  ions,  thus  :  — 

Zn(OH)2  =  Zn,  OH,  OH.  (1) 

Zinc  hydroxide  also  dissolves  in  strong  alkalies,  forming 
salts  with  these,  in  which  the  hydroxide  plays  the  role  of  an 
acid.  If  zinc  hydroxide  is  treated  with  potassium  hydroxide 
in  excess,  the  white  precipitate  dissolves,  forming  potassium 
zincate :  — 

Zn(OH)2  +  2  KOH  =  K2Zn02  +  2  H20. 

The  zinc  hydroxide,  in  the  presence  of  a  strong  base,  acts  then 
as  an  acid,  and  must  dissociate  so  as  to  yield  hydrogen  ions:  — 

Zn(OH)2  =  H,  H,  Zn02.  (2) 

The  kinds  of  ions  into  which  certain  compounds  dissociate  is 
conditioned,  then,  by  the  nature  of  the  substance  into  whose  pres- 
ence they  are  brought.  Thus,  zinc  hydroxide  in  the  presence 
of  hydrogen  ions,  dissociates  in  the  sense  of  equation  (1)  ;  in 
the  presence  of  hydroxl  ions,  in  the  sense  of  equation  (2). 

Zinc  Chloride,  ZnCL.H20. —  The  chloride  of  zinc  is  readily 
formed  by  dissolving  zinc  in  hydrochloric  acid  and  evaporating 
the  solution  to  crystallization.  When  the  salt  is  heated  to 
remove  the  water  it  loses  hydrochloric  acid,  and  a  basic  chlo- 
ride remains  behind. 

To  obtain  pure,  anhydrous,  zinc  chloride,  the  salt  with  water 
of  crystallization  must  be  heated  in  an  atmosphere  of  dry 
hydrochloric  acid  gas. 

Zinc  Sulphide,  ZnS.  —  The  sulphide  of  zinc  has  been  referred 
to  as  occurring  in  nature  as  zinc  blende,  and  as  being  one  of 
the  most  important  ores  of  zinc.  It  is  formed  as  a  white  pre- 
cipitate whenever  hydrogen  sulphide  is  conducted  into  a  dilute 
solution  of  a  soluble  zinc  salt :  — 


260  ELEMENTS   OF   INORGANIC   CHEMISTRY 

The  sulphide  of  zinc  is  soluble  to  some  extent  in  strong 
acids,  i.e.  in  the  presence  of  hydrogen  ions.  In  a  reaction  like 
the  above  where  a  strong  acid  is  liberated,  the  precipitation  of 
zinc  sulphide  is  not  complete  unless  the  solution  is  very  dilute, 
since  a  portion  of  it  is  dissolved  in  the  hydrochloric  acid 
formed  as  the  result  of  the  reaction. 

If  the  solution  of  the  zinc  salt  is  treated  with  a  solution  of 
an  alkaline  sulphide,  the  zinc  is  completely  precipitated,  since 
no  free  acid  is  formed  as  the  result  of  the  reaction :  — 

ZnCl2  -f  K2S  =  2  KC1  +  ZnS. 

Zinc  Sulphate,  ZnS04.7H20.  —  The  sulphate  of  zinc  is  ob- 
tained when  a  bar  of  pure  zinc  is  wrapped  with  a  piece  of 
platinum  wire,  dissolved  in  pure  sulphuric  acid,  and  the  solu- 
tion evaporated  to  crystallization.  The  salt  is  formed  on  a  large 
scale  by  heating  the  sulphide  in  the  presence  of  oxygen :  — 

ZnS  +  2  02  =  ZnS04. 

Like  other  sulphates  containing  a  large  number  of  molecules 
of  water  of  crystallization,  it  retains  one  molecule  to  a  much 
higher  temperature  than  the  remainder.  When  zinc  sulphate  is 
heated  slightly  above  100°  it  loses  six  molecules  of  water.  The 
last  molecule  is  retained  until  a  considerably  higher  tempera- 
ture is  reached.  On  account  of  its  color  and  composition  the 
salt  is  known  as  white  vitriol. 

Zinc  Carbonate,  ZnC03,  is  an  important  ore  of  zinc.  When 
heated  it  decomposes  into  the  oxide  and  carbon  dioxide :  — 

ZnC03=ZnO  +  C02. 

CADMIUM  (At.  Wt.  =  112.35) 

A  comparatively  rare  element  which  is  closely  allied  in  its 
-properties  to  zinc  is  cadmium.  It  usually  occurs  associated 
with  zinc  ores,  either  in  the  form  of  the  oxide  or  sulphide. 
The  oxide  of  cadmium,  like  that  of  zinc,  is  easily  reduced  by 
heated  carbon :  — 

CdO+C  =  CO  +  Cd. 


THE   MAGNESIUM  GROUP  261 


MERCURY  (At.  Wt.  =  200.0) 

Mercury  occurs  in  nature  in  the  elementary  condition,  but 
much  more  abundantly  in  the  form  of  the  sulphide,  HgS, 
cinnabar.  The  chief  localities  are  Almadem  in  Spain  and 
California.  Mercury  is  readily  obtained  by  heating  the  sul- 
phide in  contact  with  the  air.  The  metal  distils  over  and  is 
condensed,  and  the  sulphur  is  oxidized  to  sulphur  dioxide. 

Properties  of  Mercury. — Mercury  combines  with  oxygen 
only  at  elevated  temperatures,  but  combines  with  the  halogens 

at  ordinary  temperatures.     Mercury  forms  a  univalent  mercu- 

+  ++ 

rous  ion,  Hg,  and  a  bivalent  mercuric  ion,  Hg.     Both  of  these 

ions  combine  with  the  anions  of  acids,  forming  salts.  The 
former  are  known  as  mercurous,  the  latter  as  mercuric,  salts. 

Amalgams.  —  Many  of  the  metals  dissolve  readily  in  mer- 
cury, forming  amalgams,  which  are  really  solutions  of  the 
metals  in  mercury  as  a  solvent.  Such  metals  as  magnesium, 
zinc,  cadmium,  silver,  gold,  the  alkaline  earths,  the  alkalies, 
and  many  others  dissolve  without  serious  difficulty  in  mercury. 

The  amalgams  of  sodium  have  already  been  referred  to. 
The  one  containing  one  per  cent  of  sodium  is  a  liquid,  while 
double  the  amount  of  sodium  gives  a  solid  solution  in  mercury. 

The  ammonium  amalgam  formed  by  bringing  ammonium 
chloride  in  contact  with  sodium  amalgam,  has  also  been  re- 
ferred to  (p.  79). 

Mercurous  (Hg20)  and  Mercuric  (HgO)  Oxides.  —  Mercurous 
oxide  is  formed  when  a  mercurous  salt  is  treated  with  an 
alkali :  — 

2  HgCl  +  2  NaOH  =  2  NaCl  +  H20  +  Hg20. 

By  light,  heat,  or  friction,  it  is  decomposed  into  mercury  and 
oxygen.  It  is  a  black  powder  and  quite  unstable. 

Mercuric  oxide  is  formed  when  mercury  is  heated  for  a  long 
time  in  the  air.  It  is  also  obtained  as  a  red  powder  by  heating 
the  nitrate. 


262  ELEMENTS   OF  INORGANIC   CHEMISTRY 

It  is  also  formed  by  treating  a  mercuric  salt  with  sodium 
hydroxide :  — 

HgCl2  +  2  NaOH  =  2  NaCl  +  H20  +  HgO. 

Prepared  in  this  way  it  is  yellow,  but  becomes  red  on  gentle 
heating. 

Mercurous  (HgCl)  and  Mercuric  (HgCL)  Chlorides. — Mercu- 
rous  chloride  is  the  familiar  substance  calomel.  It  is  formed 
by  heating  a  mercurous  salt  with  a  soluble  chloride :  — 

Hg2S04  +  2  NaCl  =  Na2S04  +  2  HgCl. 

Mercurous  chloride  is  readily  oxidized  to  the  mercuric  salt, 
which  is  quite  poisonous. 

Mercurous  chloride  is  partly  transformed  into  mercuric 
chloride  by  the  action  of  light.  When  calomel  is  exposed  for 
a  considerable  time  to  the  action  of  light,  it  contains  the  very 
poisonous  substance  mercuric  chloride,  and  should  never  be 
used  as  a  medicine.  Calomel  which  is  to  be  used  for  such 
purposes  should  always  be  carefully  protected  from  all  agents 
which  transform  it  into  mercuric  chloride.  Mercuric  chloride, 
or  corrosive  sublimate,  is  formed  by  subliming  a  mixture  of 
sodium  chloride  and  mercuric  sulphate.  The  following  reac- 
tion takes  place :  — 

2  KaCl  +  HgS04  =  Na2S04  +  HgCl2. 

Mercuric  chloride  is  a  white,  crystalline  compound,  readily 
soluble  in  water. 

On  account  of  its  solubility  it  furnishes  a  convenient  means 

++ 
of  obtaining  mercuric  ions,  Hg,  which  are  very  poisonous  to 

most  forms  of  life.  Mercuric  chloride  is,  therefore,  an  excel- 
lent disinfectant,  and  is  extensively  used  in  surgery  for  this 
purpose. 

Mercuric  Sulphide,  HgS.  —  Only  one  compound  of  mercury 
with  sulphur  is  known,  and  this  is  the  one  in  which  the  mer- 
cury is  bivalent.  When  hydrogen  sulphide  is  brought  into  the 


THE   MAGNESIUM   GROUP  263 

presence  of  a  mercurous  salt,  a  mixture  of  black  mercuric  sul- 
phide and  mercury  is  thrown  down :  — 

2  HgCl  -f-  H2S  =  HgS  +  Hg  +  2  HC1. 

When  mercury  and  sulphur  are  rubbed  together,  the  sulphide 
is  formed.  It  is  also  formed  by  passing  hydrogen  sulphide  for 
a  time  through  a  solution  of  a  mercuric  salt :  — 

HgCl2  +  H2S  =  2  HC1  +  HgS. 

Variable  Valence.  —  We  have  studied  a  number  of  non- 
metals  which  showed  different  valence  under  different  con- 
ditions. This  is,  however,  the  first  metal  which  we  have 
encountered  in  which  two  different  valencies  clearly  manifest 

+ 

themselves.  We  have  a  well-defined  mercurous  ion,  Hg,  in 
which  the  mercury  carries  only  one  electrical  charge.  This 
forms,  with  the  anions  of  acids,  a  class  of  salts  which  have 
definite  properties.  From  these  salts  the  mercurous  mercury 
is  precipitated  by  bases  as  the  black  mercurous  oxide.  Hydro- 
chloric acid  throws  down  insoluble  mercurous  chloride.  Hy- 
drogen sulphide  precipitates  a  mixture  of  mercuric  sulphide 
and  mercury. 

The  mercuric  ion,  Hg,  has  its  own  characteristic  properties, 
forming  compounds  also  with  the  anions  of  acids.  From  these 
compounds  it  is  precipitated  by  hydrogen  sulphide  as  mercuric 
sulphide.  The  alkaline  hydroxides  precipitate  mercuric  oxide, 
while  stannous  chloride  in  excess  throws  down  metallic  mercury. 

EXPERIMENTS  WITH  MAGNESIUM,   ZINC,  CADMIUM, 
AND  MERCURY 

Experiment  162.  Magnesium  burns  in  the  Air.  —  (Mag- 
nesium ribbon.) 

Ignite  a  piece  of  magnesium  ribbon  in  a  bunsen  burner.  It 
burns  with  an  intense  white  light. 

Collect  the  magnesium  oxide  formed,  and  dissolve  in  a  large 
volume  of  water.  Test  the  reaction  of  the  solution  with  red 


264  ELEMENTS   OF  INORGANIC   CHEMISTRY 

litmus.     Add  some  of  the  solution  to  salts  of  the  heavy  metals. 
What  are  formed  ?     (See  page  257.) 

Experiment  163.     Formation  of  Magnesium  Phosphate.  - 
(Disodium  phosphate;     ammonia;     any    soluble    magnesium 
salt.) 

Add  ammonia  to  a  solution  of  clisodium  phosphate,  and  add 
some  of  this  mixture  to  any  soluble  magnesium  salt.  White 
magnesium  ammonium  phosphate  is  precipitated.  Prove  that 
ammonia  is  given  oft*  when  this  compound  is  washed,  dried, 
and  heated.  (See  page  257.) 

Experiment  164.  Zinc  Hydroxide  Soluble  in  Sodium  Hydrox- 
ide.—  (Zinc  chloride;  sodium  hydroxide.) 

Introduce  a  few  drops  of  zinc  chloride  into  a  test-tube,  and 
add  sodium  hydroxide  drop  by  drop.  At  first  zinc  hydroxide 
is  thrown  down  as  a  white  precipitate.  The  addition'  of  more 
sodium  hydroxide  dissolves  the  precipitated  zinc  hydroxide, 
and  the  solution  becomes  clear  again. 

Zinc  hydroxide  thus  differs  from  all  the  hydroxides  thus  far 
considered.  Towards  an  acid  it  acts  as  a  base,  while  towards 
a  stronger  base  it  acts  as  an  acid.  (See  page  259.) 

Experiment  165.  Precipitation  of  Zinc  Sulphide.  —  (Zinc 
chloride ;  ammonium  sulphide.) 

Add  ammonium  sulphide  to  a  solution  of  zinc  chloride. 
White  zinc  sulphide  is  precipitated.  Filter  off,  wash  with 
water,  and  treat  with  dilute  hydrochloric  acid.  Does  it  dis- 
solve ?  (See  page  259.) 

Experiment  166.  Precipitation  of  Cadmium  Sulphide.— 
(Cadmium  chloride;  hydrogen  sulphide  generator.) 

Pass  hydrogen  sulphide  through  a  cold,  dilute  solution  of 
cadmium  chloride.  Yellow  cadmium  sulphide  is  precipitated. 
Filter  off  the  precipitate,  wash  with  hot  water,  and  test  its 
solubility  in  cold,  dilute  hydrochloric  acid.  (See  page  260.) 

Experiment  167.  Formation  of  Amalgams. — (Zinc;  cad- 
mium ;  mercury.) 

Bring  a  bar  of  zinc  or  of  cadmium  into  contact  with  mercury. 
The  metal  dissolves  in  the  mercury,  which  quickly  changes  its 
appearance,  becoming  a  thick,  viscous  mass.  (See  page  260.) 

Experiment  168.  Precipitation  of  Mercury  Sulphide.  - 
(Mercuric  chloride ;  hydrogen  sulphide.) 

Test  the  solubility  of  mercurous  chloride. 

Pass  hydrogen  sulphide  into  a  solution  of  mercuric  chloride. 
Black  mercuric  sulphide  is  precipitated.  Test  the  solubility 
of  the  sulphide  in  dilute  acids.  (See  page  262.) 


THE  MAGNESIUM  GROUP  265 


PROBLEMS 

1.  Water  of  crystallization  in  50  grams  of  magnesium  sul- 
phate.    (MgS04 .  7  H20.) 

2.  What  is  the   percentage   of  magnesium    in  magnesium 
pyrophosphate  ? 

3.  How  much  zinc  is  present  in  3  grains  of  sodium  zincate  ? 
How  much  zinc  blende  is  required  to  yield  this  amount  of  zinc  ? 

4.  Weight  of  the  cadmium  sulphide  which  would  be  formed 
by  the  cadmium  in  1  gram  of  cadmium  oxide  ? 

5.  What  amount  of  mercurous  and  mercuric  chlorides  can 
be  formed  from  the  mercury  in  5  grams  of  cinnabar  ? 


CHAPTER   XXI 

ALUMINIUM  (At.  Wt.  =  27.1) 

Occurrence  and  Preparation. — Of  the  elements  of  this  group 
only  one  occurs  in  any  abundance,  and  this  is  aluminium. 
The  remainder  are  rare  substances. 

Aluminium  is  a  very  important  constituent  of  the  crust  of 
the  earth.  The  oxide  of  aluminium,  known  as  corundum,  is 
abundant,  and  the  precious  stones  ruby  and  sapphire  are  alu- 
minium oxide  colored  by  a  small  amount  of  other  substances. 

The  double  silicates  of  aluminium  and  the  alkalies  are  among 
the  most  common  minerals.  Mica  is  a  double  silicate  of  alu- 
minium and  one  of  the  alkalies,  having  the  general  composition 
MAlSi04,  in  which  M  is  a  univalent  alkali.  The  more  common 
feldspars  are  silicates  of  aluminium  and  one  of  the  alkalies 
sodium  or  potassium.  They  have  the  general  composition 
MAlSi308.  Kaolin  and  clay  are  more  or  less  pure  hydrous 
silicates  of  aluminium,  while  cryolite,  found  in  Greenland,  is  a 
double  fluoride  of  sodium  and  aluminium,  having  the  composi- 
tion 3  NaF  .  A1F3.  This  compound  is  of  importance  in  connec- 
tion with  the  preparation  of  the  element. 

Aluminium  is  prepared  to-day  by  the  electrolytic  method. 
The  compound  electrolyzed  is  aluminium  oxide.  On  account 
of  the  high  fusing-point  of  aluminium  oxide,  a  bath  of  cryolite 
is  used  and  into  the  fused  cryolite  the  aluminium  oxide  is 
introduced  as  desired.  The  cryolite  is  fused  in  iron  crucibles, 
which  are  sometimes  lined  with  carbon.  This  serves  as  the 
cathode  or  pole  which  receives  the  current  from  the  solution, 
and  upon  this  the  metal  separates;  the  oxygen  set  free  com- 
bines with  the  anode,  the  pole  from  which  the  current  enters 

266 


.  ALUMINIUM  267 

the  solution,  which  consists  of  bars  of  carbon  introduced  into 
the  fused  cryolite.  The  mass  is  kept  molten  by  the  heat  gen- 
erated by  the  current.  Since  the  introduction  of  the  electro- 
lytic method  of  preparing  aluminium,  this  metal  has  become 
quite  abundant  and  its  price  greatly  reduced. 

Properties  of  Aluminium.  —  Aluminium  is  a  metal  with 
remarkable  properties.  It  is  ductile  and  malleable  and  can  be 
readily  drawn  into  wire  or  hammered  into  thin  foil.  It  is  very 
light  for  a  metal,  having  a  specific  gravity  of  only  2.7.  It  can, 
nevertheless,  withstand  considerable  strain,  but  by  no  means  as 
great  as  was  supposed  before  it  was  prepared  on  a  large  scale. 
Its  softness  also  detracts  from  its  commercial  value. 

Chemically,  aluminium  is  fairly  resistant.  In  contact  with 
moist  air  it  becomes  covered  with  a  very  thin  layer  of  oxide, 
which  protects  the  metal  from  further  action.  It  dissolves 
readily  in  hydrochloric  acid,  while  nitric  and  sulphuric  acids 
act  only  at  elevated  temperatures. 

Aluminium  Oxide,  A1203,  and  Hydroxide,  A1(OH)3.  —  The  oxide 
of  aluminium  occurs  in  nature,  as  has  already  been  mentioned. 
Its  most  common  form  is  corundum.  Bauxite  also  occurs  in 
abundance.  Sapphire  is  a  blue  variety  used  as  a  gem,  while 
ruby  is  a  red  sapphire.  A  violet  variety  is  known  as  oriental 
amethyst,  and  a  yellow  variety  as  oriental  topaz. 

Aluminium  oxide  is  easily  prepared  by  heating  aluminium 
hydroxide  :  — 

2  A1(OH)3  =  A1203  +  3  H20. 

Aluminium  hydroxide  is  readily  prepared  by  treating  a  soluble 
aluminium  salt  with  a  base  :  — 


A1C18  4-  3  NH4OH  =  3  NH4C1  +  A1(OH),. 

It  is  also  formed  when  a  soluble  carbonate  or  sulphide  is  added 
to  a  soluble  aluminium  salt  :  — 

2  A1C1,  +  3  Na2C03  +  3  H20  =  6  NaCl  +  3  C02  +  2  A1(OH)8  ; 
2  A1C18  +  3  (NH4)2S  +6  H20  =  6  NH4C1  +  3  H2S  +  2  A1(OH)3. 


268  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Aluminium  hydroxide  is  very  slightly  soluble  in  water,  and 
is  a  very  weak  base.  Since  aluminium  hydroxide  is  a  triacid 
base,  it  must  be  capable  of  dissociating  as  follows :  — 

Al(OH),  =  A£  0~H,  OH,  OH. 

Aluminium  hydroxide  also  acts  as  an  acid  towards  strong 
bases.  When  aluminium  hydroxide  is  treated  with  a  strong 
base  like  sodium  hydroxide,  the  following  reaction  takes 
place : — 

A1(OH)8  +  3  NaOH  =  3  H20  +  Na^AlOg. 

This  compound,  known  as  sodium  ahiminate,  is  obviously  the 
sodium  salt  of  the  acid  H3A103.  The  question  which  arises 
here  is,  How  can  a  compound  which  dissociates  like  aluminium 
hydroxide,  yielding  hydroxyl  ions,  have  acid  properties  or  yield 
hydrogen  ions  ?  When  in  contact  with  a  base  or  hydroxyl  ions, 
it  dissociates  as  follows :  — 

A1(OH)3  =  H,  H2A103, 
H2A103  =  H,  HAlOa 
HA102  =  H,  A103. 

In  a  word,  it  dissociates  as  a  tribasic  acid.  We  have  seen  a 
similar  phenomenon  manifested  by  zinc,  and  other  cases  will 
appear  later. 

Cements.  —  The  alkaline  earths  also  form  aluminates.  Cal- 
cium forms  the  normal  aluminate  Ca^AlOg)^  and  also  the 
meta-aluminate  Ca(A102)2.  Since  these  salts  harden  in  contact 
with  water,  they  are  used  in  connection  with  the  preparation 
of  hydraulic  cements. 

Aluminium  Chloride,  A1C13 . 6  H20,  is  formed  when  aluminium 
hydroxide  is  treated  with  hydrochloric  acid.  From  such  a 
solution  it  can  be  obtained  in  crystalline  form  with  six 
molecules  of  water,  A1C13 .  6  H20. 

The  fluoride  of  aluminium,  as  contained  in  cryolite,  has  already 
been  referred  to  as  occurring  in  Greenland,  and  as  having  the 


ALUMINIUM  269 

composition  Na3AlF6.  It  is  obviously  the  sodium  salt  of  the 
unknown  hydrofluoraluminic  acid,  H3A1F6. 

Aluminium  Sulphate,  A12(S04)3. 18ILO,  is  formed  by  treat- 
ing aluminium  hydroxide  with  sulphuric  acid,  and  heating  the 
mixture.  On  evaporating  the  solution,  a  salt  of  the  above 
composition  separates. 

The  Alums.  —  Aluminium  sulphate  combines  with  the  sul- 
phates of  the  alkalies  to  a  remarkable  extent,  forming  a  class 
of  double  sulphates  known  as  the  alums.  These  have  the  gen- 
eral composition  MA1(SO4)2. 12H20,  in  which  M  is  an  alkali 
atom.  There  are  a  large  number  of  these  substances  ;  indeed, 
every  alkali  sulphate  forms  an  alum  with  aluminium  sul- 
phate. The  best  known  of  these  are-  potassium  alum, 
K A1(S04)2 . 12  H20,  and  ammonium  alum,  NH4A1(S04)2 . 12  H20. 
The  alums  all  crystallize  in  the  same  system,  as  octahedra  and 
cubes.  This  system,  known  as  the  regular  system,  has  all  three 
axes  of  equal  length,  and  all  make  right  angles  with  one 
another.  The  alums  are  all  isomorphous,  i.e.  form  crystals 
containing  several  of  these  substances.  When  a  crystal  of 
one  alum  is  suspended  in  a  solution  of  another  alum,  the 
second  alum  will  be  deposited  upon  it  as  upon  one  of  its  own 
crystals. 

The  term  "alum"  has  been  extended  from  the  double  sul- 
phates of  aluminium  and  the  alkalies,  to  the  double  sulphates 
of  allied  elements  and  the  alkalies.  Thus  we  have  a  series  of 
iron  alums  of  the  general  composition  MFe(S04)2 . 12  H20, 
in  which  M  is  sodium,  potassium,  rubidium,  caesium,  lithium, 
or  ammonium.  Similarly,  we  have  a  series  of  manganese 
alums,  MMn(S04)2 . 12  H20,  and  a  series  of  chromium  alums, 
MCr(S04)2 . 12  H20.  These  all  crystallize  in  the  regular  sys- 
tem, contain  twelve  molecules  of  water  of  crystallization,  and 
are  isomorphous  with  one  another  and  with  the  corresponding 
aluminium  compounds. 

Aluminium  Carbide,  A14C3.  —  Aluminium  combines  with 
carbon,  forming  the  carbide,  A14C3.  This  was  produced  by 
Moissan,  by  heating  aluminium  in  carbon  boats  in  an  electric 


270  ELEMENTS   OF  INORGANIC   CHEMISTRY 

furnace ;    also    in   preparing   aluminium    by    the    electrolytic 
method,  where  the  metal  separates  upon  carbon  electrodes. 
When   aluminium    carbide    is    decomposed    by   water,  we 

have:  A14C3  + 12  H20  =3  CH4  +  4  A1(OH)3. 

Aluminium  hydroxide  and  methane  are  formed.  Compare  this 
with  the  decomposition  of  calcium  carbide  by  water.  (Page 
242.) 

Silicates  of  Aluminium.  —  These  are  very  important  sub- 
stances. Aluminium  forms  salts  not  only  with  normal  silicic 
acid,  but  also  with  the  polysilicic  acids.  A  comparatively 
pure  form  of  aluminium  silicate  is  known  as  kaolin.  This 
substance  has  approximately  the  composition  Al4(Si04)3 .  4  H20, 
being  the  aluminium  salt  of  normal  silicic  acid. 

Clay  is  an  impure  variety  of  aluminium  silicate.  This  is 
formed  as  the  .result  of  the  weathering  of  the  rocks,  and,  con- 
sequently, many  other  substances  are  liable  to  be  present  in  it. 
The  different  colors  of  clays  are- due  to  different  impurities. 

Marl  is  clay  containing  a  large  amount  of  calcium  carbonate. 

Aluminium  silicate  readily  forms  double  silicates  with  the 
silicates  of  the  alkalies,  and  these  constitute  some  of  the  most 
important  minerals,  the  feldspars.  We  have  potassium  feld- 
spars, sodium  feldspars,  potassium  sodium  feldspars,  etc. 
These  have  the  general  composition  MAlSi308.  The  potas- 
sium feldspar  is  known  as  orthodase,  the  sodium  compound  as 
albite. 

Aluminium  silicate  forms  double  silicates  with  many  other 
elements,  especially  with  those  of  the  calcium  and  iron  group. 
Among  the  double  silicates  of  aluminium  are  such  minerals  as 
the  garnets,  mica,  zeolites,  etc. 

Applications  of  Aluminium  Silicates. — When  kaolin  or 
clay  is  mixed  with  water,  they  form  a  thick  viscous  mass, 
which  can  be  readily  moulded  or  worked  into  any  desired 
form.  When  this  mass  is  dried  and  heated,  it  becomes  very 
hard  and  then  resists  the  action  of  water.  It  is  from  this 
material  that  ordinary  brick  or  fire-brick  is  made. 


ALUMINIUM  271 

Earthenware  or  stoneware  is  made  from  a  variety  of  alumin- 
ium silicate  somewhat  purer  than  that  employed  in  the  manu- 
facture of  ordinary  fire-brick.  The  objects  are  moulded  from 
the  purer  varieties  of  clay  or  kaolin  by  mixing  with  water  to 
the  proper  consistency.  They  are  then  "  fired  "  or  "  baked  " 
by  heating  to  a  high  temperature.  The  resulting  objects  are, 
however,  very  porous,  and  would  be  comparatively  useless  in 
this  form.  They  must  be  covered  by  a  non-porous  material  — 
must  be  glazed.  There  are  different  methods  of  glazing  such 
objects.  A  mixture  of  kaolin  and  feldspar  melts  at  a  compara- 
tively low  temperature,  and  this  is  sometimes  applied  to  the 
surface  of  earthenware  after  it  has  been  "burned,"  and  the 
object  then  reheated. 

Porcelain  is  made  of  pure  aluminium  silicate,  or  pure  kaolin. 
This  is  mixed  in  definite  proportions  with  feldspar  to  lower  its 
melting-point,  and  also  with  quartz.  The  mixture  is  treated 
with  water  until  the  desired  consistency  is  reached,  and  is  then 
moulded  into  the  desired  form.  It  is  then  dried  and  "burned." 
The  glaze  consists  of  a  mixture  of  feldspar,  quartz,  and  lime. 
After  applying  the  glaze,  the  vessel  is  heated  again  to  a  very 
high  temperature.  The  many  different  varieties  of  porcelain 
owe  their  peculiar  characteristics  in  part  to  the  nature  of  the 
materials  used,  and  in  part  to  the  way  in  which  they  are 
manipulated  during  manufacture. 

Aluminium  silicate  is  also  an  important  constituent  of  many 
valuable  cements. 

In  this  same  group  belong  the  rare  elements :  scandium, 
gallium,  yttrium,  indium,  lanthanum,  and  thallium. 


EXPERIMENTS  WITH  ALUMINIUM 

Experiment  169.  Aluminium  Hydroxide  soluble  in  Sodium 
Hydroxide.  —  (Aluminium  chloride  or  an  aluminium  alum; 
sodium  hydroxide.) 

Add  sodium  hydroxide,  drop  by  drop,  to  a  solution  of  alumin- 
ium chloride  or  an  aluminium  alum.  Aluminium  hydroxide 


272  ELEMENTS   OF   INORGANIC   CHEMISTRY 

is  precipitated  as  a  white,  gelatinous  mass.  Add  more  sodium 
hydroxide,  and  the  aluminium  hydroxide  will  dissolve  again. 
Write  the  equations  ?  What  hydroxide  already  dealt  with 
does  this  suggest  ?  (See  page  267.) 

Experiment  170.  Aluminium  Hydroxide  precipitated  by  a 
Soluble  Carbonate  or  by  Ammonium  Sulphide.  —  Aluminium 
chloride  or  an  aluminium  alum ;  sodium  carbonate ;  ammo- 
nium sulphide.) 

Add  sodium  carbonate  to  a  solution  of  an  aluminium  salt. 
A  white  precipitate  resembling  aluminium  hydroxide  is 
formed.  Filter  and  wash  with  hot  water.  Add  a  little  hydro- 
chloric acid  to  the  precipitate.  It  dissolves.  Is  carbon 
dioxide  given  off?  It  is  aluminium  hydroxide.  Write  the 
equations  for  the  reactions  ? 

A  similar  precipitate  is  thrown  down  by  hydrogen  sulphide, 
since  aluminium  sulphide  is  decomposed  by  water.  Wash  ;  add 
hydrochloric  acid.  The  precipitates  dissolve.  Is  hydrogen 
sulphide  given  off  ?  The  precipitate  is  aluminium  hydroxide. 
Write  equations  ?  (See  page  267.) 

PROBLEMS 

1.  How  much  aluminium  hydroxide  can  be  prepared  from 
the  aluminium  obtained  from  2  grains  of  corundum  ?  from  5 
grams  of  cryolite  ? 

2.  Amount  of  aluminium  chloride  which  could  be  formed 
from  the  aluminium  in  8  grams  of  potassium  alum  ? 

3.  Weight  of  hydrogen  sulphide  required  to  precipitate  the 
aluminium  in  3.4  grams  of  potassium  aluminium  sulphate  ? 
How  much  aluminium  hydroxide  would  be  thrown  down  ? 


CHAPTER   XXII 

IRON,  COBALT,  NICKEL 
IRON  (At.  Wt.  =  55.9) 

WE  now  come  to  a  group  which  contains  some  of  the  most 
important  elements  technically,  as  well  as  some  of  the  most 
interesting  from  the  chemical  standpoint.  The  first  of  these, 
and  the  one  from  which  the  group  takes  its  name,  is  iron  — 
the  most  important  technically  of  all  the  elements. 

Occurrence  and  Preparation.  —  Iron  occurs  very  widely  dis- 
tributed in  nature,  but  not  in  abundance  in  the  free  state. 
This  is  due  especially  to  its  action  on  water  forming  the 
hydroxide. 

Iron  occurs  chiefly  in  the  form  of  oxides  and  sulphides.  It 
occurs  in  large  quantities  as  magnetite,  Fe304,  as  hematite, 
Fe203,  and  as  bog-iron  ore,  limonite,  and  other  hydroxides. 
Iron  also  occurs  in  large  quantities  as  pyrites,  FeS2,  and  as  the 
carbonate,  FeC03,  or  siderite. 

Iron  is  prepared  by  reduction  of  its  oxides  by  carbon  :  — 


Fe304  +  4  C  =  4  CO  +  3  Fe. 

In  preparing  iron  on  the  commercial  scale  the  blast  fur- 
nace is  employed.  This  is  shown  in  Figure  59.  The  furnace 
consists  of  an  iron  case  lined  on  the  inside  with  fire-brick,  and 
has  the  shape  shown  in  the  figure.  It  is  provided  with  pipes 
at  the  bottom  for  introducing  hot  air  under  pressure,  and  with 
pipes  at  the  top  for  carrying  off  the  gaseous  products  of  com- 
bustion. These  furnaces  are  often  quite  large,  being  as  much 
as  eighty  feet  high.  They  are  filled  with  coal  or  coke,  the  ore, 
T  273 


274 


ELEMENTS  OF  INORGANIC  CHEMISTRY 


and  a  flux,  which,  are  mixed  when  introduced  into  the  furnace. 
The  nature  of  the  flux  which  is  used  depends  upon  the  impuri- 
ties contained  in  the  ore.     If  there  is  much  silica  in  the  ore, 
limestone  is  used  as  the  flux.      The 
flux    generally    employed    is,   indeed, 
limestone. 

The  oxide  of  iron  is  reduced  to  the 
metal  by  means  of  the  highly  heated 
carbon,  and  the  carbon  monoxide 
formed  as  the  product  of  the  combus- 
tion of  the  carbon.  The  combustion 
of  the  carbon  is  effected  by  blowing 
hot  air  under  pressure  into  the  bottom 
of  the  furnaces.  Much  of  the  carbon 
monoxide  is  not  oxidized  by  the  iron 
oxide  to  carbon  dioxide,  and  escapes 
at  the  top  of  the  furnace  through  tubes 
provided  to  receive  it,  and  is  used  as 
fuel. 

The  operation  of  a  blast  furnace  is 
continuous  ;  alternate  charges  of  coke, 
ore,  and  flux  are  being  continually 
added  at  the  top  of  the  furnace,  and  the  molten  metal  and  slag 
drawn  off  at  the  bottom.  The  iron  is  run  into  moulds  made  in 
the  sand,  and  this  impure  product  is  known  as  pig-iron,  or 
cast-iron. 

Properties  of  Iron. — Pure  iron  is  light  gray  in  color,  can 
readily  be  drawn  into  wire,  and  hammered  or  rolled  into  sheets. 
At  a  bright  red  heat  it  can  be  welded,  or  one  piece  made  to 
adhere  to  another  by  simply  hammering  the  two  together. 
Iron  is  one  of  the  most  resistant  to  strain  of  all  the  metals. 
It  is  this  property,  together  with  its  great  abundance  and  the 
ease  with  which  it  can  be  prepared,  that  makes  it  the  most 
valuable  commercially  of  all  the  metals. 

When  iron  is  heated  in  the  presence  of  the  air,  it  readily 
burns,  uniting  with  oxygen  and  forming  one  of  the  oxides  of 


FIG.  59. 


IRON,    COBALT,   NICKEL  275 

iron.  Iron  acts  upon  moist  air  even  at  ordinary  temperatures, 
but  acts  slowly.  This  is  known  as  the  rusting  of  iron.  Iron 
dissolves  readily  in  dilute  acids,  liberating  hydrogen  and  com- 
bining with  the  anion  of  the  acid,  forming  the  corresponding 
salt. 

When  iron  is  dipped  into  very  strong  nitric  acid,  and  then 
into  dilute,  the  latter  is  without  action  upon  it.  The  iron  is 
then  in  the  passive  state.  This  has  recently  been  shown  to 
be  due  to  an  electrical  condition  of  the  metal,  and  not  to  the 
formation  of  a  protecting  layer  of  oxide  over  its  surface  as  was 
formerly  supposed. 

Impure  or  Commercial  Iron.  —  The  different  varieties  of  iron 
which  are  used  commercially  have  very  different  properties. 
These  are  due  to  the  different  amounts  of  impurities  in  the 
iron.  We  have  already  seen  how  pig-iron  or  cast-iron  is  made. 
Cast-iron  is  very  impure,  containing  in  addition  to  from  3  to 
4  per  cent  of  carbon,  one  or  more  per  cent  of  silicon,  besides 
phosphorus,  manganese,  sulphur,  etc.  If  there  is  considerable 
silicon  present,  and  the  cast-iron  cools  slowly,  the  carbon  sepa- 
rates largely  as  graphite,  and  gives  a  gray  cast  to  the  iron. 
This  is  known  as  gray  cast-iron.  If  the  iron  is  cooled  rapidly, 
the  carbon  remains  in  chemical  combination  with  the  iron. 
Such  iron  is  light  in  color  and  is  known  as  chilled  cast-iron. 
White  cast-iron  contains  no  graphite,  and  usually  less  silicon  or 
more  manganese  or  sulphur  than  gray  cast-iron.  Cast-iron  is 
brittle  and  readily  broken  by  a  jar,  and  is  far  less  tough  than 
pure  iron.  Cast-iron  is  not  malleable,  since  it  is  too  brittle, 
and  although  it  melts  lower  than  pure  iron,  does  not  appre- 
ciably soften  before  it  melts.  It  therefore  cannot  be  welded 
like  pure  iron. 

If  the  ore  from  which  the  iron  is  made  is  rich  in  manganese, 
the  final  product  is  also  rich  in  manganese,  and  usually  con- 
tains more  carbon  than  ordinary  cast-iron.  This  is  known  as 
spiegel  iron,  and  contains  from  10  to  15  per  cent  of  manganese, 
and  in  some  cases  even  more. 

When  most  of  the  impurities  have  been  removed  from  iron, 


276  ELEMENTS  OF   INORGANIC  CHEMISTRY 

we  have  wrought-iron.  This  still  contains  a  small  amount  of 
carbon,  the  amount,  however,  usually  being  less  than  one-half 
of  one  per  cent.  Wrought-iron  has  very  different  properties 
from  cast-iron.  It  is  very  tough,  strong,  and  malleable.  It 
melts  at  about  2000°,  but  becomes  soft  at  a  bright  red  heat,  so 
that  it  can  be  hammered,  rolled,  or  welded.  Wrought-iron, 
while  extremely  tough,  is  comparatively  soft,  and  bends  easily 
under  strain.  It  is,  therefore,  not  as  useful  as  a  form  of  iron 
which  contains  more  carbon,  and  is  known  as  steel. 

Steel  is  usually  iron  practically  free  from  all  impurities 
except  carbon,  which  is  present  to  from  0.8  to  2  per  cent. 
There  are  two  general  methods  by  which  steel  may  be  made, — 
either  by  removing  carbon  and  other  impurities  from  cast-iron, 
or  by  adding  carbon  to  wrought-iron.  The  former  process 
would  seem  to  be  the  simpler,  since  it  is  necessary  to  remove 
the  carbon  from  cast-iron  in  order  to  obtain  wrought-iron.  The 
latter  process,  however,  is  the  one  most  frequently  made  use 
of.  A  few  methods  of  making  steel  are  so  important  commer- 
cially, and  are  so  frequently  referred  to  that  they  will  be 
briefly  described. 

The  Bessemer  Converter  consists  of  a  pear-shaped  vessel  of 
malleable  iron,  lined  on  the  inside  with  refractory  material. 
The  molten  cast-iron  is  poured  into  the  converter,  and  com- 
pressed air  forced  through  the  molten  metal.  The  carbon  and 
silicon  are  completely  oxidized  by  the  oxygen  of  the  air,  and 
the  product  is  similar  in  composition  to  wrought-iron.  This  is 
kept  above  its  melting-point  by  the  heat  of  combustion  of  the 
carbon  and  silicon.  In  order  to  obtain  a  product  with  the 
desired  amount  of  carbon,  spiegel  iron  is  added  in  quantity 
sufficient  to  bring  the  percentage  of  carbon  up  to  the  desired 
amount.  The  product  is  Bessemer  steel,  which  has  found  such 
extensive  application  in  the  arts. 

The  Thomas-Gilchrist  Converter.  —  The  Bessemer  process 
of  making  steel  does  not  remove  the  phosphorus  from  the  iron. 
The  presence  of  an  appreciable  amount  of  phosphorus  so 
changes  the  properties  of  the  steel  as  to  render  it  entirely 


IRON,    COBALT,   NICKEL  277 

unfit  for  certain  purposes.  While  they  were  dependent  solely 
upon  the  Bessemer  or  similar  processes,  only  certain  iron  ores 
which  contain  only  a  small  amount  of  phosphorus,  could  be 
used  for  making  steel.  This  has  largely  been  changed,  due  to 
the  Thorn as-Gilchrist  converter.  This  is  essentially  a  Besse- 
mer converter  lined  with  burned  dolomite,  which  is  a  mixture 
of  lime  and  magnesia.  This  basic  lining,  as  it  is  termed, 
unites  with  the  phosphoric  acid  formed  by  the  oxidation  of  the 
phosphorus  by  the  oxygen  of  the  air  which  is  blown  through 
the  molten  iron,  and  forms  calcium  and  magnesium  phos- 
phates. This  material,  known  as  the  Thomas  slag,  is  exten- 
sively used  as  a  source  of  phosphoric  acid  in  commercial 
fertilizers,  having  a  composition  very  similar  to  the  "phosphate 
rock,"  which  is  so  extensively  used  as  a  fertilizer. 

Steel  is  useful  largely  because  it  can  be  made  to  assume  any 
degree  of  hardness.  When  allowed  to  cool  very  slowly,  ordi- 
nary steel  is  soft  and  resembles  wrought-iron  in  its  properties. 
When  highly  heated  and  made  to  cool  rapidly  it  becomes  very 
hard  and  brittle,  the  degree  of  hardness  depending  somewhat 
upon  the  amount  of  carbon  present.  The  process  by  which 
the  hardness  of  steel  is  regulated  is  known  as  tempering.  The 
steel  is  heated  moderately,  the  temperature  being  estimated  by 
the  color  of  the  layer  of  oxide  which  forms  on  the  bright  sur- 
face. The  steel  is  then  cooled  more  or  less  slowly. 

Ferrous  and  Ferric  Compounds.  —  Iron  forms  two  kinds  of 
ions, — the  one  carrying  two  electrical  charges  and  known  as  the 

ferrous  ion,  Fe,  and  the  other  carrying  three  electrical  charges 

+++ 
and  known  as  the  ferric  ion,  Fe.     These  are  what  have  hitherto 

been  described  as  the  ferrous  and  the  ferric  condition.  In  the 
case  of  iron  we  can  verify  the  statement  that  Faraday's  law  is 
the  base  of  chemical  valence.  If  a  given  electric  current  is 
passed  through  a  solution  of  a  ferrous,  and  a  solution  of  a 
ferric,  salt,  one  and  one-half  times  as  much  iron  will  separate 
from  the  ferrous  solution  as  from  the  ferric.  By  comparing 
the  amount  of  iron  which  separates  from  a  ferrous  solution 


278  ELEMENTS   OF  INORGANIC   CHEMISTRY 

with  the  amount  of  a  univalent  metal  separated  by  the  same 

++ 
current,  it  can  be  shown  that  the  ferrous  ion  (Fe)  is  bivalent, 

or  carries  two  charges  of  electricity.  The  ferric  ion  (Fe)  is, 
therefore,  trivalent,  or  carries  three  charges  of  electricity. 

We  shall  see  that  a  ferrous  ion  can  be  converted  into  a  ferric 
ion  by  oxidation,  as  it  is  said,  and  a  ferric  ion  converted  into 
a  ferrous  ion  by  reduction.  All  that  takes  place  when  a  fer- 
rous ion  is  converted  into  a  ferric  ion  is  the  addition  of  one 
electrical  charge,  and  the  removal  of  one  electrical  charge 
from  a  ferric  ion  converts  it  into  a  ferrous  ion.  Oxidation 
and  reduction  as  used  in  this  sense  are  simply  the  addition  or 
removal  of  electrical  energy,  and,  like  valence,  have  their  physical 
basis  in  Faraday's  law. 

Ferrous  (Fe(OH)2)  and  Ferric  (Fe(OH)3)  Hydroxides.  —  The 
two  conditions  described  above  are  exemplified  in  the  hydroxyl 
compounds  of  iron.  In  one  of  these  the  iron  holds  two 
hydroxyl  groups  in  combination  ;  in  the  other,  three.  Ferrous 
hydroxide  is  precipitated  when  an  alkali  is  added  to  a  solution 
of  a  ferrous  salt :  — 

FeCl2  +  2  NaOH  =  2  Nad  +  Fe(OH)2. 

Ferric  hydroxide  is  formed  when  an  alkali  is  added  to  a 
solution  of  a  ferric  salt :  — 

FeCl3  4-  3  NaOH  =  3  NaCl  +  Fe(OH)3. 

Iron  forms  two  well-known  oxides,  Fe203,  or  hematite,  and 
Fe304,  or  magnetite. 

Ferrous  (FeCl2)  and  Ferric  (FeCl3)  Chlorides.  —  Ferrous 
chloride  is  obtained  when  an  excess  of  iron  is  treated  with 
hydrochloric  acid.  It  forms  crystals  containing  four  mole- 
cules of  water,  FeCl2 . 4  H20. 

Ferric  chloride  is  formed  by  passing  chlorine  over  iron.  It 
is  also  formed  by  passing  chlorine  into  a  solution  of  ferrous 
chloride. 

Sulphides  of  Iron.  —  When  iron  filings  and  sulphur  are  heated 
together,  the  two  combine  and  form  ferrous  sulphide,  FeS.  It  is 


IRON,   COBALT,   NICKEL  279 

also  formed  by  the  action  of  ammonium  sulphide  on  a  ferrous 
salt.  Treated  with  acids,  hydrogen  sulphide  is  liberated.  Two 
other  sulphides  of  iron  exist. 

Ferrous  Sulphate,  FeS04.  7  H20.  —  Ferrous  sulphate,  also 
called  iron  vitriol  on  account  of  its  composition,  or  green 
vitriol  on  account  of  its  color,  is  the  most  important  ferrous 
compound.  It  is  formed  by  the  action  of  sulphuric  acid  on  iron 
or  iron  sulphide.  It  is  made  commercially  by  allowing  ferrous 
sulphide  to  take  up  oxygen  from  the  air :  — 

FeS  +  2  02  =  FeS04. 

The  salt  is  then  extracted  with  water.  Ferrous  sulphate 
forms  light  green  crystals,  and  is  extensively  used  in  dyeing, 
in  pharmacy,  and  as  a  disinfectant. 

Ferric  Sulphate,  Fe2(S04)3,  is  formed  by  dissolving  ferric 
oxide  or  hydroxide  in  sulphuric  acid :  — 

2  Fe(OH)3  +  3  H2S04  =  Fe2(S04)3  +  6  H2O. 

It  is  also  formed  by  the  addition  of  a  half-equivalent  of  sul- 
phuric acid  to  ferrous  sulphate,  in  the  presence  of  an  oxidizing 
agent  like  nitric  acid.  Ferric  sulphate  forms  double  sulphates 
with  the  alkali  sulphates,  which,  in  composition  and  crystalline 
form,  resemble  the  aluminium  alums,  and  are  termed  iron  alums. 
Potassium  Ferrocyanide,  K4Fe(CN)6. — Although  iron  does 
not  combine  directly  with  the  cyanogen  ion  and  form  cyanides, 
it  forms  double  cyanides  which  are  beautifully  crystallized  com- 
pounds. When  potassium  cyanide  is  allowed  to  act  upon  iron 
in  the  presence  of  water,  the  following  reaction  takes  place :  — 

Fe  +  6  KCN  +  2  H2O  =  2  KOH  +  H2  +  K4Fe(CN)6. 

Potassium  ferrocyanide  is  usually  formed  by  heating  nitrog- 
enous matter  with  iron  filings  and  potash.  When  the  mass  is 
digested  with  water  and  the  solution  evaporated,  beautiful  yellow 
crystals  are  formed  having  the  composition  K4Fe(CN)6.3  H20. 
This  is  potassium  ferrocyanide,  known  commercially  as  yellow 
prussiate  of  potash. 


280  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Potassium  Ferricyanide,  K3Fe(OT)6,  is  formed  by  the  action 
of  oxidizing  agents  on  potassium  ferrocyanide.  If  chlorine  is 
passed  into  a  solution  of  potassium  ferrocyanide,  potassium 
f erricyanide  is  formed :  — 

2  K4Fe(CN)6  +  C12  =  2  KC1  +  2  K3Fe(CN)6. 

The  compound  K3Fe(CN)6  is  known  also  as  red  prussiate  of 
potash,  from  the  deep  red  color  of  its  crystals. 

Change  in  Color  with  Change  in  Electrical  Charge.  —  An  ion 
having  the  same  chemical  composition  does  not  always  have 
the  same  color.  Take  the  ion  Fe(CN)6;  in  potassium  ferro- 
cyanide it  is  yellow  and  gives  the  yellow  color  to  a  solution  of 
this  salt.  The  ion  in  this  case  is  formed  by  the  dissociation 
of  the  salt  K4Fe(CN)6 :  — 

K4Fe(CN)6  =  K,  K,  K,  K,  Fe(CN)6. 

= 

The  ion  Fe(CN)6  carries  four  negative  charges.  The  ion 
Fe(CN)6,  obtained  by  the  dissociation  of  potassium  f  erricya- 
nide, is  red.  The  compound  K3Fe(CISr)6  dissociates  as  fol- 
lows:—  +  +  +  = 

K3Fe(CN)6  =  K,  K,  K,  Fe(CN)«. 

The  ion  Fe(CN)6,  in  this  case,  carries  three  negative  charges, 
and  the  difference  of  one  charge  changes  the  color  of  the  ion 
from  yellow  to  red. 

To  take  a  simpler  example:    The  iron  ion  in  the  ferrous 

++ 
condition,  Fe,  is  green,  as  is  seen  in  solutions  of  ferrous  salts ; 

while  the  iron  ion  in  the  ferric  condition,  Fe,  is  yellow,  as  is 
seen  in  solutions  of  ferric  salts.  A  large  number  of  examples 
of  changes  in  the  color  of  ions  with  change  in  the  electrical 
charge  which  they  carry,  might  be  given. 

COBALT  (At.  Wt  =  59.0) 

The  mineral  smaltite,  which  is  the  arsenide  of  the  composi- 
tion CoAs2,  and  the  mineral  cobaltite,  which  has  the  composition 


IRON,    COBALT,    NICKEL  281 

CoAsS,  are  important.  The  element  is  prepared  by  reducing 
the  oxide  either  with  highly  heated  carbon,  or  with  hydrogen. 
Cobalt  resembles  iron  in  many  respects. 

Cobaltous  and  Cobaltic  Compounds.  — Cobalt,  like  iron,  forms 

++ 
two  kinds  of  ions,  —  the  cobaltous  ion  Co,  and  the  cobaltic  ion 

4-4-+ 

Co.  Unlike  iron,  however,  the  cobaltous  ion  is  the  more  stable 
condition,  while  the  ferric  condition  is  the  more  stable  for 
the  iron  ion.  The  cobaltous  ion  combines  readily  with  the 
anions  of  acids,  forming  solutions  of  cobalt  salts. 

Oxides  and  Hydroxides  of  Cobalt.  —  Several  oxides  of  cobalt 
are  known.  Cobaltous  oxide,  CoO,  is  formed  when  cobaltous 
carbonate  or  hydroxide  is  heated.  It  is  a  greenish  powder, 
easily  reducible  to  the  metal.  Cobaltic  oxide,  or  cobalt  sesqui- 
oxide,  Co203,  is  formed  by  gently  heating  the  nitrate. 

When  a  cobaltous  salt  is  treated  with  an  alkali,  cobaltous 
hydroxide  is  formed  :  — 

CoCl2  +  2  KOH  =  2  KC1  +  Co(OH)2. 

Cobaltous  Salts.  —  When  metallic  cobalt  is  heated  in  chlorine, 
cobalt  chloride,  CoCl2,  is  formed.  This  crystallizes  from  an 
aqueous  solution  with  six  molecules  of  water  —  CoCl2 .  6  H20. 
The  solution  of  the  salt  is  red.  When  the  water  is  removed, 
the  salt  is  deep  blue  in  color,  due  to  the  driving  back  of  the 
ions  into  molecules,  which  are  blue.  Cobalt  chloride  has, 
therefore,  been  used  as  sympathetic  ink.  When  a  solution  of 
the  salt  is  used  for  writing  on  paper,  the  writing  is  practically 
colorless,  due  to  the  nearly  colorless  nature  of  an  aqueous 
solution  of  cobaltous  chloride. 

When  the  paper  is  warmed  the  blue  color  appears,  and  the 
writing  becomes  plainly  legible.  When  the  blue  material  is 
allowed  to  stand  in  contact  with  the  air,  it  takes  up  moisture, 
becoming  again  invisible. 


282  ELEMENTS   OF   INORGANIC   CHEMISTRY 

NICKEL   (At.  Wt.  =  58.7) 

An  element  which  resembles  cobalt  in  many  respects  is 
nickel.  It  occurs  in  nature  chiefly  in  combination  with  arse- 
nic, NiAs,  as  niccolite.  The  silicate  occurs  in  abundance  and  is 
known  as  garnierite. 

Nickel  is  prepared  by  reducing  the  oxide  with  carbon  at  a 
high  temperature,  and  by  reducing  the  oxide  in  a  stream  of 
hydrogen. 

On  account  of  its  resistance  to  oxidation,  nickel  is  extensively 
used  to  cover  metals  which  are  more  readily  oxidized,  such  as 
iron,  etc.  The  nickel  is  deposited  upon  the  iron  electrolytically. 
This  method  of  covering  one  metal  with  another  is  known  as 
electroplating. 

Nickel  forms  valuable  alloys  with  a  number  of  the  metals. 
German  silver  is  an  alloy  of  nickel  with  zinc  and  copper.  Alloys 
of  nickel  and  copper  are  used  as  coins ;  our  so-called  nickel 
containing  75  per  cent  copper  and  25  per  cent  nickel. 

Compounds  of  Nickel.  —  While  there  are  a  few  compounds 

known  in  which  nickel  plays  the  role  of  a  trivalent  element,  it 

++ 
is  almost  always  present  as  the  bivalent  ion  Ni.     The  oxide  of 

nickel,  NiO,  is  formed  as  a  black  powder  when  the  hydroxide 
is  heated  in  a  limited  supply  of  air.  When  there  is  an  abun- 
dant supply  of  oxygen,  the  sesquioxide,  Ni203,  is  formed. 

The  green  hydroxide,  Ni(OH)2,  is  formed  when  a  solution  of 
a  nickel  salt  is  treated  with  a  solution  of  a  hydroxide :  — 

NiCl2  +  2  KOH  =  2  KC1  +  Ni(OH)2. 

EXPERIMENTS   WITH  IRON,  COBALT,   AND  NICKEL 

Experiment  171.  Preparation  of  Ferrous  Chloride. —  (Iron 
wire  or  iron  filings  ;  hydrochloric  acid.) 

Treat  5  grams  of  iron  wire  or  iron  filings  with  15  cc.  of 
dilute  hydrochloric  acid.  Smell  the  odor  of  the  escaping  gases. 

Treat  a  little  of  the  ferrous  chloride  with  sodium  hydroxide. 
Ferrous  hydroxide  is  precipitated.  What  is  its  color  ?  As  it 
stands  in  contact  with  the  air  it  passes  over  into  ferric  hydrox- 


IRON,   COBALT,   NICKEL  283 

ide.  Note  the  changes  in  color  which  it  undergoes  as  it  passes 
from  the  practically  white  ferrous  hydroxide  to  the  reddish 
brown  ferric  hydroxide. 

Treat  a  little  ferrous  chloride  with  a  little  concentrated 
nitric  acid  and  boil  the  solution.  The  ferrous  ion  is  gradually 
transformed  into  the  ferric  ion.  Note  the  change  in  the  color 
of  the  solution.  This  same  transformation  is  effected  by  pass- 
ing chlorine  through  the  solution.  (See  page  277.) 

Experiment  172.  Properties  of  Ferrous  Salts. — (Ferrous 
sulphate ;  potassium  ferricyanide.) 

Treat  a  solution  of  potassium  ferricyanide  with  a  few  drops 
of  a  solution  of  ferrous  sulphate.  A  deep  blue  precipitate  is 
formed, —  TurnbuWs  blue.  Perform  the  same  experiment,  using 
potassium  ferrocyanide.  A  light  blue  color  is  produced. 

Treat  a  solution  of  ammonium  sulphocyanate  with  a  solution 
of  a  ferrous  salt.  If  there  is  no  ferric  salt  present  in  the  ferrous 
salt,  the  solution  of  the  sulphocyanate  will  not  change  color. 
This  is  secured  by  boiling  the  solution  just  before  using  with  a 
little  iron  or  zinc  and  hydrochloric  acid.  (See  page  277.) 

Experiment  173.  Separation  of  Ferric  Chloride;  Properties 
of  Ferric  Salts.  —  (Iron  wire  or  iron  filings ;  hydrochloric  acid ; 
nitric  acid.) 

Prepare  ferric  chloride  as  at  the  end  of  Experiment  171. 
Dissolve  the  iron  in  hydrochloric  acid.  Add  nitric  acid  and 
boil  the  solution.  Kepeat  the  addition  of  nitric  acid  and  the 
boiling  two  or  three  times  to  transform  all  of  the  ferrous  com- 
pound into  the  ferric  condition. 

Add  to  a  little  of  the  ferric  chloride  a  solution  of  potassium 
ferrocyanide.  A  characteristic  blue  precipitate  is  formed,— 
Berlin  or  Prussian  blue. 

Add  a  few  drops  of  ferric  chloride  to  a  solution  of  potassium 
sulphocyanate.  A  characteristic  blood-red  color  appears. 

Add  a  few  drops  of  a  ferric  salt  to  a  solution  of  potassium 
ferricyanide;  no  blue  color  is  produced.  The  solution  turns 
brown.  (See  page  277.) 

Experiment  174.  Detection  of  a  Ferrous  and  a  Ferric  Salt 
when  Both  are  Present.  —  (Ferrous  sulphate ;  ferric  chloride.) 

Mix  a  gram  of  ferrous  sulphate  with  a  gram  of  ferric 
chloride  and  test  the  mixture  for  ferrous  and  ferric  iron. 

Add  to  a  few  drops  of  the  solution  a  solution  of  potassium 
ferricyanide.  What  does  the  blue  color  show  ? 

Add  a  few  drops  of  the  solution  to  a  solution  of  potassium 
ferrocyanide.  What  does  the  blue  color  indicate  ? 


284  ELEMENTS   OF  INORGANIC   CHEMISTRY 

To  a  solution  of  potassium  sulphocyanate  add  a  drop  of  the 
solution  of  the  mixed  iron  salts.  The  red  color  indicates  the 
presence  of  iron  in  what  condition  ? 

State  clearly  and  concisely  the  reactions  for  a  ferrous  salt  ? 
for  a  ferric  salt  ?  (See  page  278.) 

Experiment  175.  Precipitation  of  Iron  as  the  Sulphide.  — 
(Ferrous  sulphate ;  ferric  chloride ;  hydrogen  sulphide  gener- 
ator; ammonium  sulphide.) 

Into  a  solution  of  ferrous  sulphate,  and  of  ferric  chloride, 
conduct  hydrogen  sulphide.  What  takes  place  ? 

To  similar  solutions  of  the  same  substances  add  ammonium 
sulphide.  What  is  formed  in  each  case  ? 

Test  the  solubility  of  the  precipitates  in  dilute  acids  ? 

Explain  why  hydrogen  sulphide  does  not  precipitate  all  of 
the  iron  from  solutions  of  ferrous  salts  ?  (See  page  278.) 

Experiment  176.     Sympathetic  Ink.  —  (Cobalt  chloride.) 

Prepare  a  fairly  concentrated  solution  of  cobalt  chloride. 
Write  upon  paper  with  this  solution.  Allow  the  paper  to  dry. 
Can  anything  be  seen  ?  Warm  the  paper  carefully  over  the 
flame  or  on  the  air-bath.  The  characters  come  out  in  blue. 
Explain  ?  (See  page  281.) 

Experiment  177.  Action  of  Hydrogen  Sulphide  on  Cobalt 
Salts. — (Cobalt  chloride;  hydrogen  sulphide  generator.) 

Pass  hydrogen  sulphide  into  a  solution  of  cobalt  chloride. 
What  is  precipitated  ?  Is  it  soluble  in  dilute  acids  ?  Add 
ammonia.  What  is  formed  ?  Explain  ? 

Experiment  178.  Action  of  Hydrogen  Sulphide  on  Nickel 
Salts.  —  (Nickel  chloride ;  hydrogen  sulphide  generator.) 

Pass  hydrogen  sulphide  into  a  solution  of  a  nickel  salt. 
What  is  formed  ?  Test  its  solubility  in  dilute  acids.  Add 
ammonia.  What  is  precipitated?  Explain  the  difference 
between  the  action  of  hydrogen  sulphide  and  of  ammonium 
sulphide  ? 

PROBLEMS 

1.  How  much  ferrous  and  ferric  chlorides  can  be  prepared 
from  the  iron  contained  in  3  grams  of  hematite  ?  in  3  grams 
of  magnetite  ? 

2.  Amounts  of  potassium  ferrocyanide  and  potassium  ferri- 
cyanide  which  contain  10  grams  of  iron  ? 

3.  Percentage   of    cobalt   in   the    compound    CoCl2 . 6  H20. 
Amount  of  water  in  5  grams  of  this  substance? 

4.  How  much  niccolite  would  be  required  to  yield  the  nickel 
necessary  to  form  100  grams  of  nickel  sulphide  ? 


CHAPTER  XXIII 

MANGANESE,  CHROMIUM,  MOLYBDENUM,  TUNGSTEN, 

URANIUM 

MANGANESE  (At.  Wt.  =  55.0) 

WE  now  come  to  an  element  which  probably  forms  as  large 
a  variety  of  compounds  as  any  element  known.  This  is  due 
to  the  many  degrees  of  valence  which  manganese  can  manifest. 

Occurrence,  Preparation,  and  Properties  of  Manganese. — 
Manganese  occurs  in  nature  in  small  quantities  in  the  free  con- 
dition, but  usually  as  one  of  its  oxides.  The  chief  source  of 
manganese  is  the  oxide  Mn02,  which  is  the  mineral  pyrolusite. 

Manganese  is  prepared  by  heating  the  oxides  with  carbon  in 
an  electric  furnace,  also  by  electrolysis  of  the  fused  chloride, 
but  more  conveniently  by  mixing  the  oxide  with  finely  divided 
aluminium  and  igniting  the  mixture.  This  is  one  of  Gold- 
Schmidts  mixtures,  the  aluminium  taking  the  oxygen,  setting 
free  the  manganese. 

Oxides  of  Manganese.  —  Manganese  forms  no  less  than  seven 
compounds  with  oxygen.  The  one  containing  the  smallest 
amount  of  oxygen  is  manganous  oxide,  MnO.  This  is  formed 
by  reducing  the  higher  oxides  in  a  stream  of  hydrogen,  and 
by  heating  manganous  hydroxide.  Manganese  sesquioxide, 
Mn203,  occurs  in  nature  as  braunite.  Manganous-manganic 
oxide,  Mn3O4,  is  formed  when  the  other  oxides  of  manganese 
are  heated  in  the  air.  Manganese  dioxide,  Mn02,  occurs  in 
nature  as  pyrolusite,  and  is  the  most  important  ore  of  manga- 
nese. There  exists  a  trioxide  of  manganese,  Mn03,  and  also  a 
septoxide,  Mn207.  There  also  exists  a  tetroxide  of  manganese, 
Mn04. 

285 


286  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Arranging  these  oxides  in  the  order  of  increasing  amount  of 
oxygen,  we  have :  — 

MnO  Mn03 

Mn304  Mn207 

Mn203  Mn04 

Mn02 

++ 
Manganous   Salts.  —  The    manganous    ion,   Mn,    combines 

with  the  anions  of  acids,  forming  salts  which  are  usually 
beautifully  crystallized  compounds.  Manganous  chloride, 
MnCl2 .  4  H20,  is  formed  by  the  action  of  hydrochloric  acid  on 
manganese  dioxide :  — 

Mn02  +  4  HC1  =  2  H20  +  C12  +  MnCl2. 

+++ 
Manganic  Compounds.  —  The  manganic  ion,  Mn,  can  unite 

with  the  anions  of  acids  and  form  salts.  The  manganic  salts 
are,  however,  not  as  numerous  as  the  manganous,  and  in 
general  not  as  stable,  being  strongly  hydrolyzed  by  water. 

Quadrivalent  Manganese.  —  A  few  compounds  are  known  in 
which  tetravalent  manganese  apparently  exists.  This  is  the 
case  with  the  oxide,  Mn02,  and  the  hydroxide.  The  most 
important  of  these  substances  is  manganese  dioxide,  which, 
as  we  have  seen,  formerly  found  extensive  application  in  the 
preparation  of  chlorine,  and  to-day  is  largely  used  in  the  arts 
as  an  oxidizing  agent. 

It  is  also  used  in  the  construction  of  one  of  the  most  efficient 
forms  of  primary  cells,  the  Ledanche  cell 

Valence  and  Properties  of  Manganese.  —  It  should  be  noted 
that  as  the  valence  of  manganese  increases,  its  basic  nature 
rapidly  diminishes.  Bivalent  manganese  is  distinctly  basic, 
forming  stable  salts  with  the  anions  of  acids.  Trivalent  man- 
ganese is  very  weakly  basic,  its  salts  being  strongly  hydro- 
lyzed by  water.  Tetravalent  manganese  is  scarcely  basic  at 
all,  its  compound  with  such  a  strong  acid  as  hydrochloric 
being  so  unstable  that  its  very  existence  is  doubtful. 

When  we  pass  to  manganese  with  higher  valence,  not  only 
has  all  the  basic  nature  been  lost,  but  we  find  acid  properties 


MANGANESE  287 

beginning  to  manifest  themselves,  and  the  highest  oxidation 
product  of  manganese  is  a  strong  acid.  These  acid  compounds 
of  manganese  we  shall  now  study. 

Manganic  Acid,  H2Mn04. — We  have  already  studied  com- 
pounds of  manganese  in  which  this  element  has  appeared  in  the 
capacity  of  a  bivalent,  trivalent,  and  quadrivalent  ion.  Pen- 
tavalent  manganese  is  not  known,  but  hexavalent  manganese  is 
well  known,  manifesting  itself  in  salts  of  the  compound,  man- 
ganic acid,  —  the  analogue  of  sulphuric  acid.  These  are  formed, 
as  we  would  expect,  by  strongly  oxidizing  manganese  in  the 
presence  of  bases.  Potassium  manganate.  K2Mn04,  is  formed 
by  fusing  potassium  hydroxide  with  manganese  dioxide  in  the 
presence  of  the  oxygen  of  the  air,  or  better,  an  oxidizing  agent 
such  as  potassium  chlorate.  The  manganese  is  oxidized  from 
the  quadrivalent  to  the  sexivalent  condition,  the  potassium 
chlorate  being  reduced  to  potassium  chloride  :  — 

3  Mn02  +  KC103  +  6  KOH  =  KC1  +  3  H20  +  3  K2Mn04. 

This  mass  forms  a  green  solution,  from  which  green  crystals  of 
potassium  manganate,  K2Mn04,  separate. 

The  following  transformation  is  effected  by  carbon  dioxide 
and  consequently  takes  place  slowly  when  a  manganate  is 
exposed  to  the  air:  — 

3  K2Mn04  +  2  C02  =  Mn02  +  2  K2C03  +  2  KMn04. 

The  change  of  color  from  the  green  manganate,  through 
blue  and  purple  to  the  purplish  red  permanganate,  is  very 
striking. 

Permanganic  Acid,  KMn04.  —  The  highest  oxidation  product 
of  manganese  containing  hydrogen  and  oxygen  is  permanganic 
acid,  HMn04,  —  the  analogue  of  perchloric  acid.  The  best 
method  of  preparing  the  acid  consists  in  the  electrolysis  of  the 
potassium  salt  of  this  acid.  This  method,  which  was  devised 
by  Morse,  is  carried  out  as  follows :  Two  unglazed  porcelain 
cups  containing  the  one  water  to  which  a  little  alkali  is  added, 


288  ELEMENTS   OF   INORGANIC   CHEMISTRY 

and  the  other  water  to  which  a  little  permanganic  acid  is  added 
if  available  to  make  the  water  conducting,  are  immersed  in  a 
beaker  containing  a  solution  of  potassium  permanganate.  The 
platinum  electrodes  are  inserted  the  one  in  each  cup,  the  cathode 
in  the  cup  containing  the  alkali.  The  current  is  passed,  when 
the  potassium  ions  move  toward  the  cathode,  give  up  their 
charge,  decompose  water,  arid  liberate  hydrogen.  The  alkali 
formed  around  the  cathode  is  easily  siphoned  off  from  time  to 

time.  The  permanganic  ions,  Mn04,  move  to  the  anode,  decom- 
pose water,  forming  permanganic  acid  and  liberate  oxygen. 
The  permanganic  acid  collects  in  the  cup  around  the  anode, 
and  after  the  current  has  been  passed  for  a  sufficient  time,  can 
be  obtained  in  perfectly  pure  condition  and  in  any  quantity 
desired.  Permanganic  acid  is  a  very  strong  acid. 

Potassium  permanganate,  KMn04,  is  readily  obtained  by  pass- 
ing carbon  dioxide  through  a  solution  of  potassium  manganate, 
as  already  described,  or  by  treating  a  carbonate  with  perman- 
ganic acid.  Its  solution  has  exactly  the  same  color  as  per- 
manganic acid,  —  purplish  red.  It  crystallizes  in  beautiful 
purple  crystals. 

Potassium  permanganate  is  characterized  chiefly  by  its  oxi- 
dizing power. 

CHROMIUM  (At.  Wt.  =  52.1) 

Chromium  forms  a  number  of  series  of  compounds,  and 
many  of  these  are  closely  related  to  iron  and  manganese.  It 
occurs  in  nature  largely  as  chrome  iron  ore,  which  is  iron 
chromite,  having  the  composition  Fe(Cr02)2. 

Chromium  is  prepared  most  conveniently  by  heating  the 
oxide  with  finely  divided  aluminium,  according  to  the  method 
of  Goldschmidt. 

Oxides  of  Chromium.  —  Chromium,  like  manganese,  forms  a 
number  of  oxides.  Chromous  oxide,  CrO,  is  formed  by  reduc- 
ing the  higher  oxides.  Chromic  oxide  or  chromium  sesquioxide, 
Cr203,  is  formed  by  heating  the  trioxide.  Chromium  trioxide, 


CHROMIUM  289 

Cr03,  is  formed  by  adding  concentrated  sulphuric  acid  to  potas- 
sium dichromate  iu  very  concentrated  solution:  — 

K,Cr  A  +  H2S04  =  K2S04  -f  H20  +  2  Cr03. 
It  is  a  dark  red,  beautifully  crystalline  substance,  character- 
ized by  its  tremendous  oxidizing  power.      When  brought  in 
contact  with  organic  compounds,  these  are  oxidized  or  burned 
up,  as  we  say. 

Valence  and  Properties  of  Chromium  Ions.  — It  is  obvious, 
from  the  composition  of  the  oxides  and  hydroxides  of  chro- 
mium, that  this  element  can  exist  in  various  conditions  of  va- 

++ 

lence.     Bivalent  chromium  ions,  Cr,  are  distinctly  basic.     The 

+++ 
bivalent  ions,  however,  readily  pass  into  trivalent  ions,  Cr, 

which  are  very  weakly  basic  towards  strong  acids,  and  are  acidic 
towards  certain  bases.  This  is  strictly  analogous  to  the  ions  of 
iron  and  manganese.  Those  of  lower  valence,  or  with  the 
smaller  electrical  charge,  are  basic  j  but  as  the  valence  increases, 
or  as  the  amount  of  electrical  energy  which  they  carry  increases, 
the  basic  property  becomes  less  and  less,  and  acidic  properties 
begin  to  manifest  themselves.  When  the  valence  of  the  chro- 
mium ion  reaches  six,  as  in  the  compound  H2CrO4,  we  have  a 
very  strong  acid,  chromic  acid,  and,  similarly,  when  it  reaches 
seven,  in  perchromic  acid,  HCr04.  This  is  analogous  to  iron, 
and  especially  to  manganese,  where  the  sexivalent  ion  is  acidic 
as  in  manganic  acid,  and  the  septivalent  ion  strongly  acidic  as 
in  permanganic  acid. 

Chromic  Acid,  H2Cr04.  —  In  composition  this  acid  resembles 
sulphuric  acid,  manganic  acid,  and  the  like. 

The  compound  H2Cr04  is,  however,  not  known.  When  its 
salts  are  treated  with  sulphuric  acid,  the  anhydride  Cr03  is 
obtained,  which  is  chromic  acid  minus  water:  — 

H2Cr04-H2O  =  Cr03. 

In  this  compound  the  chromium  is  obviously  sexivalent,  and 
with  its  high  valence  the  strongly  acid  properties  begin  to 
come  out. 
u 


290  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Chromates.  —  Salts  of  this  acid  are  formed  when  ferrous 
chromite  is  heated  on  the  air  in  the  presence  of  an  alkali. 
The  chromium  is  oxidized  to  the  chromate,  which  is  soluble,  and 
the  iron  to  ferric  oxide,  which  is  insoluble  in  water.  The  mixture 
is  treated  with  water,  when  the  chromate  dissolves.  If  caustic 
potash  is  used,  the  resulting  compound  is  potassium  chromate. 
This  is  deep  yellow  in  color,  due  to  the  color  of  the  chromic 
acid  ion  Cr04.  It  forms  crystals  which  are  isomorphous  with 
potassium  sulphate. 

Bichromates.  —  When  a  chromate  like  potassium  chromate 
is  treated  with  an  acid,  the  color  of  the  solution  changes  from 
yellow  to  red.  The  reaction  in  the  case  of  potassium  chromate 
may  be  represented  as  follows :  — 

2  K2Cr04  +  H2S04  =  H20  +  K2S04  +  K2O207. 

Potassium  dichromate,  K2Cr207,  crystallizes  from  the  solution 
in  beautiful  red  crystals,  which  often  grow  to  unusual  size  and 
are  of  unusual  geometrical  perfection.  Potassium  dichromate 
is  a  powerful  oxidizing  agent.  It  readily  gives  up  oxygen,  the 
chromium  being  reduced  to  the  chromic  condition. 

When  potassium  dichromate  is  treated  with  caustic  potash, 
potassium  chromate  is  formed :  — 

K2O207  +  2  KOH  =  2  K2Cr04  +  H20. 

This  is  made  evident  by  the  change  in  the  color  of  the  solu- 
tion from  red  to  yellow.  It  will  be  observed  that  the  valence 
of  the  chromium  in  potassium  dichromate  is  the  same  as  in 
potassium  chromate.  The  change  from  the  former  to  the  latter 
by  the  addition  of  an  alkali,  and  the  reverse  change  by  the 
addition  of  an  acid,  therefore,  involve  neither  oxidation  nor 
reduction. 

URANIUM  (At.  Wt.  =  238.5) 

The  three  comparatively  rare  elements  molybdenum,  tungsten, 
and  uranium  also  belong  in  this  group. 


MANGANESE,  CHROMIUM  291 

Uranium  is  of  special  interest  on  account  of  its  power  of 
emitting  a  certain  remarkable  kind  of  radiation. 

Uranium  Radiation.  —  Compounds  of  uranium  have  the  prop- 
erty of  emitting  an  invisible  radiation,  which  traverses  many 
substances  impervious  to  light,  such  as  black  paper,  thin  sheets 
of  many  metals,  such  as  aluminium,  copper,  etc.  This  property 
is  possessed  by  metallic  uranium  to  from  three  to  four  times 
the  extent  that  it  is  manifested  by  the  salts  of  this  metal. 

This  is  entirely  different  from  the  phosphorescence  shown  by 
salts  of  uranium,  since  the  latter  disappears  very  quickly,  while 
the  power  of  emitting  this  invisible  radiation  persists  for 
years. 

If  a  piece  of  uranium  or  of  one  of  its  salts  is  placed  above  a 
photographic  plate  covered  with  black  paper  or  aluminium 
leaf,  and  various  substances  interposed  between  the  uranium 
and  the  plate,  after  several  hours  radiographs  are  obtained 
upon  the  plate. 

Other  radioactive  substances  are:  radium,  polonium,  and 
thorium. 

EXPERIMENTS  WITH  MANGANESE  ,AND  CHROMIUM 

Experiment  179.  Reaction  of  Manganese  Salts.  —  (Soluble 
salts  of  manganese  ;  ammonia ;  sodium  hydroxide ;  hydrogen 
sulphide  generator;  ammonium  sulphide.) 

Treat  solutions  of  manganese  salts  with  ammonia.  What  is 
precipitated  ?  Add  sodium  hydroxide  to  solutions  of  man- 
ganese salts.  What  is  precipitated  ?  Is  it  soluble  in  an  excess 
of  sodium  hydroxide?  How  does  manganese  differ  in  this 
respect  from  aluminium  ? 

Pass  hydrogen  sulphide  into  solutions  of  manganese  salts. 
Is  any  precipitate  formed  ? 

Add  ammonium  sulphide  to  solutions  of  manganese  salts. 
What  is  precipitated  ?  What  is  the  color  of  the  precipitate  ? 
Is  it  soluble  in  dilute  acids  ?  Why  is  a  manganese  salt  not 
precipitated  by  hydrogen  sulphide  ? 

Experiment  180.  Preparation  of  Potassium  Manganate. — 
(Porcelain  crucible ;  manganese  dioxide ;  potassium  hydroxide  ; 
potassium  chlorate.) 


292  ELEMENTS  OF  INORGANIC   CHEMISTRY 

Mix  3  grams  of  finely  powdered  manganese  dioxide  with  3 
grams  of  potassium  hydroxide  and  add  2  grams  of  potassium 
chlorate.  Heat  the  mixture.  It  will  turn  green,  due  to  the 
formation  of  potassium  inanganate.  Write  the  equation  for 
the  reaction  ? 

Dissolve  the  green  mass  in  a  little  water,  and  pass  carbon 
dioxide  through  the  solution.  It  will  gradually  become  purple 
in  color.  This  is  due  to  the  oxidation  of  the  inanganate  to  the 
permanganate.  This  same  transformation  is  effected  by  allow- 
ing the  crystals  of  the  manganate  to  stand  exposed  to  the  air. 
(See  page  287.) 

Experiment  181.  Color  of  Potassium  Permanganate.— 
(Potassium  permanganate;  dilute  sulphuric  acid.) 

When  potassium  permanganate  dissolves  in  water,  it  dissoci- 
ates thus :  — 

KMn04  =  K,  Mn04. 

We  naturally  ask  to  which  ion  the  beautiful  color  of  the 
solution  is  due.  It  may  be  due  to  both. 

That  it  is  not  due  to  the  potassium  ion  is  shown  by  the  fact 
that  solutions  of  such  potassium  salts  as  the  chloride,  nitrate, 
sulphate,  etc.,  are  colorless.  These  yield  in  solution  large 
numbers  of  potassium  ions,  and  yet  tiieir  solutions  are  color- 
less. The  color  must  be  due  to  the  Mn04  ion. 

Treat  a  solution  of  potassium  permanganate  with  dilute  sul- 
phuric acid.  Permanganic  acid,  HMn04,  is  formed.  This  dis- 
sociates as  follows :  — 

HMn04  =  H,  Mn04. 

Explain  why  permanganic  acid  in  solution  has  the  same 
color  as  potassium  permanganate.  (See  page  287.) 

Experiment  182.  Oxidizing  Power  of  Potassium  Permanga- 
nate. —  (Potassium  permanganate  ;  oxalic  acid  ;  potassium  sul- 
phite; sulphuric  acid;  ferrous  sulphate.) 

To  a  dilute,  warm  solution  of  oxalic  acid  to  which  sulphuric 
acid  has  been  added,  add  a  solution  of  potassium  permanganate, 
drop  by  drop.  The  color  of  the  permanganate  will  disappear 
as  long  as  there  is  any  oxalic  acid  present.  The  oxalic  acid  is 
oxidized  to  water  and  carbon  dioxide,  and  the  latter  is  seen  to 
escape  from  the  solution.  The  manganese  is  reduced  to  the 
bivalent  condition,  and  forms  manganous  sulphate. 


MANGANESE,  CHROMIUM  293 

Add  sulphuric  acid  to  a  solution  of  potassium  sulphite.  Sul- 
phurous acid  is  set  free  in  the  solution.  Add  potassium  per- 
manganate, drop  by  drop.  The  sulphurous  acid  is  oxidized  to 
sulphuric  acid,  and  the  manganese  reduced  as  in  the  above 
experiment. 

Add  sulphuric  acid  to  a  dilute  solution  of  ferrous  sulphate, 
and  then  add  potassium  permanganate,  drop  by  drop.  The 
color  instantly  disappears  on  coming  in  contact  with  the  ferrous 
salt.  The  ferrous  sulphate  is  oxidized  to  ferric  sulphate,  and 
the  permanganate  reduced. 

Add  a  drop  of  potassium  permanganate  to  the  end  of  your 
finger.  The  organic  matter  is  oxidized,  and  the  stain  of  oxide  of 
manganese  shows  that  the  permanganate  has  been  reduced. 
The  stain  can  be  readily  removed  by  adding  a  drop  of  sul- 
phurous acid.  (See  page  288.) 

Experiment  183.  Formation  of  Chromium  Trioxide.  —  (Po- 
tassium dichromate ;  concentrated  sulphuric  acid.) 

Fill  a  test-tube  one-third  full  of  potassium  dichromate  and 
cover  the  crystals  with  water.  Warm  carefully.  When  no 
more  of  the  salt  will  dissolve,  pour  the  solution  into  a  second 
test-tube  and  add  carefully,  drop  at  a  time,  concentrated  sul- 
phuric acid.  Red  crystals  of  chromium  trioxide  will  separate. 
These  must  not  be  allowed  to  come  in  contact  with  the  hand. 
Filter  the  solution  through  glass  wool  (not  filter-paper).  By 
means  of  a  porcelain  spatula  remove  a  few  of  the  crystals  and 
lay  them  upon  filter-paper.  What  occurs  ?  Treat  a  few  of 
the  crystals  with  a  little  hydrochloric  acid  and  warm.  Smell 
the  gas  given  off.  Explain  ?  (See  page  289.) 

Experiment  184.  Chromates  and  Bichromates.  —  (Potassium 
chromate ;  potassium  dichromate  ;  silver  nitrate ;  lead  nitrate ; 
sulphuric  acid ;  hydrochloric  acid.) 

Dissolve  25  grams  of  potassium  chromate  in  a  little  water 
and  add  sulphuric  acid.  The  solution  changes  from  yellow  to 
red.  Evaporate  the  solution,  and  beautiful  red  crystals  of 
potassium  dichromate  separate.  Write  the  equation  for  the 
above  reaction. 

Treat  a  solution  of  potassium  dichromate  with  caustic  potash, 
and  the  color  changes  from  red  to  yellow.  Explain  ?  Write 
the  equation  for  the  above  reaction  ? 

Add  a  solution  of  potassium  chromate  to  a  solution  of  a  salt 
of  a  heavy  metal.  The  chromate  of  the  metal  is  precipitated. 
The  chromate  of  the  metal  is  also  precipitated  when  a  solution 
of  a  dichromate  is  used  instead  of  the  chromate.  Write  equa- 


294  ELEMENTS   OF  INORGANIC   CHEMISTRY 

tions  for  the  reaction  between  potassium  chromate  and  lead 
nitrate ;  for  the  reaction  between  potassium  chromate  and  sil- 
ver nitrate. 

Note  the  color  of  the  precipitates  in  the  two  cases  ?  (See 
page  290.) 

Experiment  185.  Oxidizing  Power  of  the  Chromates  and 
Bichromates.  —  (Potassium  chromate;  potassium  dichromate; 
concentrated  hydrochloric  acid.) 

Treat  a  solution  of  potassium  chromate  and  of  dichromate 
in  separate  test-tubes  with  concentrated  hydrochloric  acid. 
Heat  the  solutions.  What  evidence  have  you  that  the  hydro- 
chloric acid  is  oxidized  ?  Is  there  any  evidence  that  the 
chromate  and  dichromate  are  reduced  ? 

To  a  solution  of  potassium  dichromate  in  a  test-tube  add 
concentrated  hydrochloric  acid  and  a  little  alcohol.  The  solu- 
tion quickly  turns  green,  showing  that  the  chromium  has  been 
reduced  to  the  trivalent  or  chromic  condition.  Smell  the  odor 
of  the  escaping  vapor.  This  is  the  odor  of  aldehyde  formed 
by  the  oxidation  of  the  alcohol.  (See  page  290.) 

Experiment  186.  Action  of  Reagents  on  Chromic  Chromium. 
—  (Potassium  chromium  alum  ;  sodium  hydroxide ;  potassium 
carbonate  ;  ammonium  sulphide.) 

Potassium  chrome  alum  or  potassium  chromium  sulphate, 
KCr(S04)2,  is  the  analogue  of  potassium  aluminium  sulphate, 
KA1(S04)2.  It  is  obvious  that  in  the  above  compound  the 
chromium  is  in  the  chromic  or  trivalent  condition. 

When  potassium  chromium  sulphate  is  treated  with  sodium 
hydroxide,  chromium  hydroxide  is  precipitated.  Write  the 
equation  ? 

Addition  of  more  sodium  hydroxide  dissolves  the  precipi- 
tate. Write  the  equation  ?  Is  this  similar  to  the  behavior  of 
aluminium  hydroxide  ? 

Boil  the  solution ;  the  chromium  hydroxide  is  reprecipitated. 
Is  this  analogous  to  aluminium  ? 

Treat  another  portion  of  the  solution  of  chrome  alum  with 
sodium  carbonate.  Chromium  hydroxide  is  precipitated. 
Write  the  equation?  It  is  exactly  analogous  to  aluminiumv 
Treat  still  another  portion  of  the  solution  of  the  chromium" 
alum  with  ammonium  sulphide.  Chromium  hydroxide  is  pre- 
cipitated, since  the  sulphide  is  decomposed  by  water.  Write 
the  equation  from  the  analogy  to  aluminium  ? 


PROBLEMS  295 


PROBLEMS 

1.  How  much  potassium  manganate  and  how  much  potas- 
sium permanganate  contain  the  same  amount  of  manganese  as 
2  grams  of  manganite  ? 

2.  Amount  of  potassium  chromate  and  of  potassium  dichro- 
mate  obtainable  from  10  grams  of  the  oxide,  Cr203? 


CHAPTER   XXIV 

COPPER,  SILVER,  GOLD 
COPPER  (At.  Wt.  =  63.6) 

THERE  still  remain  three  elements  in  the  first  group  of  the 
Periodic  System  which  have  not  thus  far  been  studied.  These 
are  copper,  silver,  and  gold. 

Occurrence  and  Preparation  of  Copper.  —  Copper  occurs  in 
considerable  quantity  in  the  free  condition.  It  occurs  in 
large  quantities  as  cuprous  oxide,  CuO,  or  cuprite  ;  cupric  oxide ; 
azurite  and  malachite  the  blue  and  green  basic  carbonates ;  and 
chalcopyrite. 

Copper  is  prepared  from  the  oxides  very  simply  by  heating 
with  charcoal. 

Copper  is  finally  purified  by  means  of  electrolysis.  The  impure 
copper  is  moulded  into  the  form  of  large,  thick  plates,  known 
as  the  anode  plates.  These  are  suspended  in  a  large  bath  of 
copper  sulphate,  and  are  made  the  anodes  by  connecting  them 
with  the  positive  pole  of  a  dynamo.  Between  these  plates  are 
alternately  suspended  thin  sheets  of  pure  copper,  which  are  the 
cathodes,  and  these  are  connected  with  the  negative  pole  of  the 
dynamo.  When  the  current  is  passed,  copper  is  deposited  upon 
each  of  the  cathodes,  and  dissolves  from  each  of  the  anodes. 
The  action  of  the  current  is  really  to  carry  the  copper  from  t}ie 
anode  to  the  cathode  opposite  to  it,  and  deposit  it  upon  the 
cathode. 

Under  these  conditions  the  impurities  are  not  deposited  with 
the  copper,  but  either  remain  in  solution  or  are  deposited,  in  the 
form  of  a  viscous  mass  on  the  bottom  of  the  copper  sulphate  bath. 

Properties  of  Copper.  —  Copper  differs  in  color  from  all  other 
metals,  being  a  peculiar  shade  of  red  known  as  copper-red. 

296 


COPPER,   SILVER,    GOLD  297 

Copper  is  quite  resistant  to  chemical  reagents.  In  contact  with 
inoist  air  it  becomes  covered  with  a  green  basic  carbonate. 
When  heated  in  the  air,  it  forms  the  oxide. 

Copper  is  easily  attacked  by  sulphur  compounds,  forming 
the  sulphide. 

On  account  of  its  physical  properties,  copper  is  one  of  the 
most  valuable  of  the  metals.  It  can  be  readily  hammered  into 
thin  sheets  or  drawn  into  wire,  and  is  very  strong. 

Next  to  silver,  copper  is  the  best  conductor  of  electricity,  and 
i&  extensively  used  in  this  capacity  in  connection  with  teleg- 
raphy and  telephony,  and  especially  in  connection  with  electric 
lighting  and  electrotraction,  where  large  amounts  of  electrical 
energy  must  be  transported.  This  is  one  of  the  most  important 
uses  of  the  element  copper. 

Alloys  of  Copper.  —  With  the  metals  copper  forms  a  number 
of  alloys  which  are  very  valuable.  One  of  the  best  known 
is  brass,  which  is  an  alloy  of  copper  and  zinc,  containing  gen- 
erally about  twice  as  much  copper  as  zinc ;  but  this  varies 
greatly  from  one  specimen  to  another.  German  silver  or 
argentan,  as  we  have  seen,  is  an  alloy  of  copper,  nickel,  and 
zinc.  Copper  also  forms  alloys  with  nickel  and  silver.  These 
are  frequently  used  for  coins.  The  silver  coins  usually  contain 
about  ten  per  cent  of  copper. 

Among  the  best-known  alloys  of  copper  are  the  bronzes. 
The  ordinary  bronzes  are  alloys  of  copper  and  tin,  containing 
from  10  to  30  per  cent  of  tin.  Among  the  alloys  of  copper  and 
tin  are  bell  metal  and  gun  metal. 

Oxides  of  Copper.  —  Copper  forms  two  compounds  with  oxy- 
gen, —  cuprous  oxide,  Cu20,  and  cupric  oxide,  CuO.  These  are 
types  of  the  two  classes  of  copper  compounds,  —  the  cuprous 
compounds,  in  which  copper  is  univalent,  and  the  cupric  com- 
pounds, in  which  the  copper  is  bivalent.  Copper,  therefore, 

forms  two  kinds  of  ions,  —  the  cuprous  ion,  Cu,  and  the  cupric 

++ 
ion,  Cu.     Of  these  the  cupric  condition,  in  which  the  copper 

carries  two  electrical  charges,  is  the  more  stable. 


298  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Copper  Sulphate,  CuS04  .  5  H20.  —  This  is  the  best  known  of 
all  the  compounds  of  copper,  and  one  of  the  best-known  sub- 
stances. Copper  sulphate,  or  blue  vitriol,  as  it  is  called,  is 
formed  when  sulphuric  acid  acts  on  metallic  copper.  The 
hydrogen  is  not  set  free  but  acts  on  more  sulphuric  acid, 
reducing  it  to  sulphur  dioxide  :  — 

H2S04=CuS04 


Copper  sulphate  crystallizes  from  the  solution  in  the  form  of 
beautiful  blue  crystals,  containing  five  molecules  of  water  — 
CuS04  .  5  H20.  When  the  sulphate  is  heated,  it  loses  four 
molecules  of  water  at  100°,  but  the  fifth  is  retained  until  200° 
is  reached. 

Precipitation  of  Copper  by  Zinc  and  Iron.  —  When  a  bar  of 
zinc  or  iron  is  immersed  in  a  solution  of  a  copper  salt,  copper 
is  precipitated  upon  the  zinc  or  iron  and  the  zinc  or  iron  dis- 
solves. WThen  zinc  replaces  copper  from  its  salts,  the  zinc  atom 
takes  the  charge  from  the  copper  ion,  becoming  itself  an  ion, 
while  the  copper  is  converted  into  an  atom  :  — 

Zn  +  Cu,  S54  =  Cu  +  Zn,  S04. 


SILVER   (At.  Wt.  =  107.93) 

We  now  come  to  the  so-called  noble  metals  or  precious 
metals.  The  well-known  elements  silver  and  gold  will  now 
be  studied. 

Silver  is  not  among  the  rare  elements.  It  occurs  in  nature 
in  considerable  abundance  and  in  a  number  of  compounds.  An 
important  ore  is  the  sulphide  Ag2S,  argentite. 

Preparation  of  Silver.  —  A  large  number  of  methods  have 
been  devised  and  used  for  obtaining  pure  silver.  The  method 
employed  depends  upon  the  nature  of  the  silver  ore  which  is 
being  used. 

When  silver  is  set  free  along  with  many  other  substances,  it 


COPPER,   SILVER,    GOLD  299 

is  frequently  dissolved  in  mercury  and  the  mercury  then  dis- 
tilled off.  This  is  known  as  the  amalgamation  process. 

Lead  ores,  especially  galena,  usually  contain  silver,  and  sil- 
ver is  frequently  obtained  mixed  with  lead.  The  solution  of 
silver  in  lead  is  concentrated  by  allowing  it  to  crystallize. 
Pure  lead  separates  at  first,  and  the  remaining  solution  becomes 
richer  and  richer  in  silver.  A  concentration  of  the  silver  in 
the  lead  is  finally  attained,  where  the  crystals  which  separate 
have  the  same  concentration  of  silver  as  the  remaining  solution. 

When  this  stage  is  reached,  it  is  impossible  to  effect  further 
separation  by  crystallization.  The  above  process  of  fractional 
crystallization  can  be  continued  until  the  solution  contains 
about  one  per  cent  of  silver. 

To  effect  further  separation  the  lead  solution  of  silver  is 
heated  on  the  air.  The  lead  is  oxidized  to  litharge,  and  allowed 
to  flow  away  or  be  absorbed  by  the  porous  walls  of  the  cupel 
in  which  the  oxidation  takes  place.  This  process  is  known  as 
cupellation. 

Properties  of  Silver.  —  Silver  is  a  white  metal  and  has  a  high 
lustre.  It  is  not  as  hard  as  copper,  and  of  all  the  metals  is 
the  best  conductor  of  heat  and  electricity.  It  can  be  easily 
drawn  into  wire  or  hammered  into  thin  foil. 

Silver  is  not  easily  attacked  by  chemical  reagents.  It  is  not 
attacked  by  the  strongest  alkalies  even  when  hot,  nor  by  dilute 
acids,  with  the  exception  of  nitric  acid. 

Silvering.  —  Silver,  as  we  have  seen,  is  quite  resistant  to 
ordinary  chemical  agents.  It  is  consequently  used  for  mak- 
ing utensils  and  objects  of  ornament.  These  are,  however,  ex- 
pensive, and  silver-plated  wares  are  much  used  in  their  stead. 
These  consist  of  brass,  copper,  or  other  metallic  objects  covered 
completely  with  metallic  silver.  They,  therefore,  'have  the 
properties  of  silver  objects. 

Silver-plating  finds  extensive  application.  The  silver  is  de- 
posited by  a  number  of  methods ;  the  silver  is  reduced  directly 
upon  the  object  to  be  plated,  or  it  is  applied  mechanically  and 
pressed  upon  the  object  while  hot.  Another  method  which 


300  ELEMENTS   OF   INORGANIC    CHEMISTRY 

is  extensively  used  is  the  electrolytic.  Silver  is  deposited 
electrolytically  from  a  solution  of  the  cyanide  dissolved  in 
potassium  cyanide. 

Silver  is  now  being  extensively  deposited  upon  glass  in  the 
construction  of  mirrors. 

Oxides  and  Hydroxide  of  Silver.  —  Silver  forms  three  com- 
pounds with  oxygen,  —  the  suboxide,  Ag40,  the  normal  oxide, 
Ag20,  and  the  superoxide,  AgO.  It,  however,  forms  only  one 
hydroxide,  AgOH,  which  is  stable  only  at  very  low  tempera- 
tures. 

The  silver  ion  is  univalent,  combining  with  the  anions  of 
acids  and  forming  salts  of  the  general  type,  AgCl,  AgX03, 
Ag2S04,  etc. 

The  silver  ion  is  especially  a  reagent  for  the  halogen  ions. 
It  combines  with  them,  forming  stable,  insoluble  compounds, 
which  we  shall  now  study. 

Silver  Chloride,  AgCl,  is  formed  whenever  a  silver  ion  comes 
in  contact  with  a  chlorine  ion  :  — 

A+g  +  Cl  =  AgCl. 

It  is  a  white  precipitate,  which  quickly  darkens  when  exposed 
to  the  light. 

Silver  chloride  is  practically  insoluble  in  water,  and  is  con- 
sequently used  to  determine  quantitatively  both  silver  and 
chlorine.  It  is  readily  soluble  in  aqueous  ammonia,  and  is 
thus  distinguished  from  the  bromide  and  iodide. 

Silver  Bromide,  AgBr.  —  Silver  bromide  is  precipitated  when 
silver  ions  come  in  contact  with  bromine  ions  :  — 


Silver  bromide  is  white  with  a  slightly  yellowish  tint;  is 
soluble  with  difficulty  in  ammonia  and  almost  insoluble  in 
water.  Silver  bromide  is  even  more  sensitive  to  light  than 
silver  chloride,  and  upon  this  fact  is  based  its  use  in  photog- 
raphy. 


COPPER,   SILVER,   GOLD  301 

Photography.  —  The  science  of  photography  is  based  almost 
exclusively  upon  the  action  of \light  on  silver  bromide.  The 
sensitive  film  is  prepared  by  adding  ammonium  bromide  to 
gelatine,  and  then  adding  silver  nitrate  in  the  dark.  The  fol- 
lowing reaction  takes  place :  — 

NH4Br  +  AgN03  =  NH4N03  +  AgBr. 

The  silver  bromide  is  distributed  through  the  gelatine  in  a  very 
fine  state  of  division.  The  mass  is  then  poured  upon  the  surface 
of  glass  plates,  to  which  it  adheres  in  the  form  of  a  thin  film. 

The  plate  containing  the  film  is  then  exposed  to  the  action  of 
the  light  from  the  object  which  it  is  desired  to  photograph. 
The  time  of  the  exposure  depends  upon  the  intensity  of  the 
light,  and  the  sensitiveness  of  the  film.  The  action  of  the 
light  is  probably  to  reduce  the  silver  bromide  to  a  sub-bromide 
of  silver,  although  this  is  not  proved. 

The  exposed  plate  is  now  treated  with  a  developer,  which 
consists  of  some  reducing  agent,  such  as  pyrogallic  acid,  ferrous 
sulphate,  etc.  The  object  of  the  developer  is  to  reduce  the  silver 
bromide  depositing  metallic  silver,  and  the  whole  science  of 
photography  depends  upon  the  fact  that  the  silver  bromide 
which  is  most  strongly  illuminated  is  most  readily  reduced 
by  the  developer.  Where  the  object  was  brightest,  the  plate 
is  covered  with  a  deeper  film  of  metallic  silver,  which  becomes 
less  and  less  dense  as  the  illumination  is  less  and  less.  The 
result  is  a  photograph  of  the  object  with  the  light  parts  dark 
and  the  dark  parts  light.  This  is  the  so-called  "negative." 
The  negative  thus  obtained  still  contains  unreduced  silver 
bromide,  since  the  portion  of  the  salt  which  was  not  exposed 
to  the  light  is  not  reduced  by  the  developer.  If  the  negative  is 
exposed  to  the  light  in  this  condition,  the  silver  bromide  would 
be  acted  upon,  and  the  original  picture  would  be  destroyed. 
To  avoid  this  the  negative  must  be  fixed,  i.e.  treated  with  a 
solution  of  a  substance  which  will  dissolve  the  unreduced  silver 
bromide.  The  fixing  agent  usually  employed  is  a  solution  of 
sodium  thiosulphate,  Na2S203,  known  technically  as  hyposul- 


302  ELEMENTS   OF   INORGANIC   CHEMISTRY 


)  or  even  as  hypo.  This  acts  upon  the  silver  bromide, 
forming  a  double  salt,  which  is  quite  soluble  in  water  and  is 
easily  removed  when  the  plate  is  washed  with  running  water. 
The  negative  is  now  "  fixed,"  and  ready  to  be  used  in  mak- 
ing prints  or  positives.  A  "positive  picture,"  or  a  photo- 
graph proper,  is  obtained  by  placing  the  negative  above  paper 
covered  with  the  sensitive  film,  and  exposing  it  to  light.  The 
dark  parts  of  the  negative  cut  off  the  light  and  appear  bright 
on  the  positive  picture,  and,  conversely,  the  light  parts  appear 
dark,  since  much  light  passes  through  and  acts  upon  the  sensi- 
tive film  upon  the  paper.  The  positive,  therefore,  represents 
the  lights  and  shades  in  the  order  in  which  they  occur  in  the 
object,  and  is  a  true  picture  in  metallic  silver  of  that  object. 

In  some  cases  the  print  is  immersed  in  a  bath  containing  a 
gold  or  platinum  salt,  when  the  silver  precipitates  the  gold  or 
platinum,  itself  passing  into  solution.  Such  photographs  have, 
then,  the  soft  brown  color  of  finely  divided  gold,  or  the  harder 
steel-gray  tint  of  finely  divided  platinum. 

Silver  Iodide,  Agl,  is  formed  whenever  silver  and  iodine  ions 
come  in  contact.  It  is  a  yellow  solid,  insoluble  in  water  and  in 
ammonia.  Silver  iodide,  like  the  bromide  and  chloride,  is  sen- 
sitive to  light,  and  was  formerly  used  in  connection  with 
photography.  Indeed,  the  earliest  method  of  preparing 
photographs,  devised  by  Daguerre,  made  use  of  silver  iodide. 

GOLD  (At.  Wt.  =  197.25) 

The  element  gold  is  one  of  the  "  noble  metals,"  and  is  fre- 
quently classed  with  platinum  and  allied  elements.  On  the 
whole,  however,  it  seems  best  to  study  gold  in  connection  with 
copper  and  silver. 

Gold  occurs  in  nature  chiefly  in  the  uncombined  condition,  in 
the  form  of  nuggets  or  grains  in  quartzite  rocks  or  in  sands.  It 
also  occurs  combined  with  the  element  tellurium,  as  the  telluride. 
When  it  occurs  native,  it  is  by  no  means  pure,  containing  silver, 
copper,  etc. 


COPPER,    SILVER,   GOLD  303 

The  Metallurgy  of  Gold.  — Gold  occurs  usually  in  very  small 
quantities,  and  widely  distributed  through  a  large  mass  of  rock 
or  sand.  It  must  be  obtained  free  from  large  quantities  of 
foreign  substances.  This  is  accomplished  by  placer  mining.  In 
placer  mining  the  earth  or  sand  is  washed  with  water,  the  light 
materials  being  carried  away  and  the  gold  left  behind  with  the 
heavier  substances.  In  hydraulic  mining  the  gold-bearing  earth 
and  sands  are  washed  down  from  the  hills  by  water  under 
pressure,  and  the  heavy  gold  collected  by  dissolving  it  in  mer- 
cury. The  gold  is  obtained  from  the  amalgam  by  distilling  off 
the  mercury. 

If  certain  impurities  are  present,  the  cyanide  process  is  used. 
This  consists  in  treating  the  gold  ore  with  potassium  cyanide, 
in  which  finely  divided  gold  readily  dissolves.  The  gold  is 
precipitated  from  the  cyanide  solution  by  means  of  metallic 
zinc,  or  electrolytically,  and  then  subjected  to  cupellation. 
Gold  thus  obtained  is  impure  and  must  be  purified.  The  sil- 
ver can  be  removed  by  dissolving  it  out  in  nitric  or  concen- 
trated sulphuric  acid.  If  the  amount  of  gold  in  the  alloy 
exceeds  25  per  cent,  this  does  not  work  satisfactorily.  In 
such  cases  the  alloy  is  fused  with  enough  silver  to  dilute  the 
gold  to  not  more  than  one-fourth.  The  process  is  therefore 
called  quartation. 

Properties  of  Gold.  —  Gold  is  a  soft,  yellow  solid,  melting  at 
1064°,  and  forming  a  greenish  liquid.  It  has  a  very  high 
specific  gravity,  — 19.3.  Gold  is  extremely  malleable,  and  can 
be  hammered  into  leaves  not  more  than  two-millionths  of  a 
millimetre  in  thickness.  Such  gold  leaf  is  translucent  and  has 
a  green  color. 

Gold  is  very  resistant  chemically,  being  attacked  by  com- 
paratively few  substances. 

The  solution  of  gold  in  chlorine  water  is  of  special  interest. 
When  the  molecules  of  gold  come  in  contact  with  the  mole- 
cules of  chlorine,  the  former  become  cations  and  the  latter 
anions :  — 

Au  +  Cl  +  Cl  +  Cl  =  Au+,  Cl,  Cl,  Cl. 


304  ELEMENTS   OF  INORGANIC  CHEMISTRY 

Gold  forms  alloys  with,  a  number  of  the  metals.  The  best 
known  and  the  most  important  are  the  alloys  with  copper  and 
silver.  Pure  gold  is  too  soft  for  use  either  as  coin  or  as  orna- 
mental objects.  To  make  it  harder  and  more  durable,  copper 
is  added.  This  gives  to  the  gold  a  deep  red  color.  The  alloy 
containing  10  per  cent  of  copper  is  frequently  used.  The 
purity  of  the  gold  is  expressed  in  carats,  pure  gold  being  24 
carats.  The  number  of  carats  means  the  number  of  parts  of 
gold  in  24  parts  of  the  alloy.  Thus,  18-carat  gold  means  an 
alloy  containing  18  parts  gold  and  6  parts  copper. 

The  alloy  of  gold  and  silver  is  extensively  used  instead  of 
pure  gold,  being  more  resistant  to  abrasion  and  more  durable. 

Gold-plating.  —  Metal  objects  are  covered  with  gold  in  the 
same  manner  and  for  the  same  purpose  that  they  are  covered 
with  silver.  Gold-plating  has  been  effected  by  a  number  of 
methods,  but  these  have  practically  all  given  way  to  the  elec- 
trical. The  object  to  be  electroplated  with  gold  is  made  the 
cathode,  and  a  piece  of  pure  gold  the  anode,  the  double  cyanide 
of  gold  and  potassium  being  the  electrolyte. 

Oxides  and  Hydroxides  of  Gold.  —  Gold  forms  the  two  oxides 
Au20  and  Au203,  which  are  typical  of  the  univalent  and  triva- 
lent  compounds  of  gold.  It  also  forms  the  corresponding 
hydroxides,  Au(OH)  and  Au(OH)3.  Although  these  com- 
pounds are  weak  bases,  combining  with  the  anions  of  certain 
acids  and  forming  salts,  the  auric  hydroxide  also  has  acid 
properties.  + 

Salts  of  Gold.  —  Gold   forms   the  aurous  ion  Au,  and  the 

auric  ion  Au.  One  of  these  carries  one  electrical  charge,  or  is 
univalent,  and  the  other  carries  three  electrical  charges,  or  is 
trivalent.  These  ions  can  form  salts  with  the  anions  of  certain 
acids. 

EXPERIMENTS  WITH   COPPER   AND   SILVER 

Experiment  187.     Action  of  Reagents  on  Salts  of  Copper.  — 

(Copper  sulphate;  sodium  hydroxide;  ammonium  sulphide; 
hydrogen  sulphide.) 


COPPER,   SILVER,   GOLD  305 

To  a  cold,  dilute  solution  of  copper  sulphate  add  sodium 
hydroxide.  Cupric  hydroxide  is  precipitated  with  a  bluish 
color.  Warm  the  solution ;  the  color  gradually  changes  to 
black.  This  is  due  to  the  loss  of  water  and  the  formation  of 
cupric  oxide. 

Pass  hydrogen  sulphide  through  a  solution  of  copper  sul- 
phate. What  occurs  ? 

Add  ammonium  sulphide  to  a  solution  of  copper  sulphate. 
What  is  precipitated  ?  Write  equation  ? 

Test  the  solubility  of  copper  sulphide  in  dilute  acids?  (See 
page  298.) 

Experiment  188.  Copper  precipitated  from  its  Salts  by 
Certain  Metals.  —  (Copper  sulphate ;  bar  of  zinc.) 

File  the  surface  of  a  bar  of  zinc  to  clean  it,  and  plunge  it  into 
a  solution  of  copper  sulphate.  Copper  is  deposited  upon  the 
zinc,  and  zinc  passes  into  solution.  The  reaction  that  takes 
place  is  the  following  :  — 

Cu,  S04  +  Zn  =  Zu,  S04  -f  Cu. 

Zinc  takes  the  positive  charge  from  the  copper,  itself  becoming 
an  ion,  while  the  copper  having  lost  its  charge  becomes  an 
atom,  and  is  precipitated  in  the  metallic  condition.  (See  page 
298.) 

Experiment  189.  Copper  precipitates  Certain  Metals  from 
their  Salts.  —  (Strip  of  copper ;  mercuric  chloride  or  nitrate.) 

Dip  a  strip  of  copper  into  a  solution  of  mercuric  chloride  or 
mercuric  nitrate.  Mercury  is  precipitated  upon  the  copper, 
and  some  of  the  copper  dissolves.  The  reaction  that  takes 
place  is :  — 

Hg,  CI,  CI  +  Cu  =  Cu,  C~l,  Ci  +  Hg. 

Copper  takes  the  charge  from  the  mercury,  becoming  an  ion, 
while  the  mercury,  having  lost  its  charge,  becomes  an  atom. 

Experiment  190.  Action  of  Reagents  on  Silver  Nitrate.  — 
(Silver  nitrate;  sodium  hydroxide;  potassium  chloride,  bro- 
mide, and  iodide;  ammonia;  hydrogen  sulphide.) 

Treat  silver  nitrate  with  sodium  hydroxide.  What  is 
formed?  Is  it  soluble  in  an  excess  of  sodium  hydroxide  ? 

Treat  silver  nitrate  with  potassium  chloride.  What  is 
formed  ?  Write  the  equation  ? 


306  ELEMENTS   OF   INORGANIC   CHEMISTRY 

Treat  silver  nitrate  with  potassium  bromide.  What  is  pre- 
cipitated ?  Write  equation  ? 

Treat  silver  nitrate  with  potassium  iodide.  What  is  the 
precipitate  ?  Equation  ? 

Test  the  solubility  of  silver  chloride,  bromide,  and  iodide  in 
ammonia  ? 

Pass  hydrogen  sulphide  through  a  solution  of  silver  nitrate. 
What  is  precipitated  ?  Is  it  soluble  in  dilute  acids  ? 

Experiment  191.  Action  of  Light  on  Silver  Chloride  and 
Silver  Bromide.  —  (Silver  nitrate ;  potassium  chloride  ;  potas- 
sium bromide.) 

Precipitate  some  silver  chloride.  Preserve  a  part  of  it  in 
the  dark  and  expose  the  remainder  to  the  light  for  a  few 
hours.  What  difference  do  you  notice  in  the  two  samples  ? 

Precipitate  silver  bromide  and  preserve  a  part  in  the  dark, 
exposing  the  remainder  to  the  light.  What  visible  effect  is 
produced  by  the  light  ?  (See  page  301.) 

Experiment  192.  Silver  precipitated  by  Mercury.  —  (Silver 
nitrate;  mercury.) 

Dissolve  3  grams  of  silver  nitrate  in  100  cc.  water.  Place 
the  solution  in  a  narrow  cylinder  and  add  one  or  two  drops 
of  mercury.  Silver  will  separate,  and  mercury  will  pass  into 
solution.  The  reaction  is  expressed  thus:  — 

Ag,  N03  +  Hg  =  Ag  +  Hg,  N03. 

The  mercury  removes  the  charge  from  the  silver,  becoming  an 
ion.     The  silver,  having  lost  its  charge,  becomes  an  atom. 


PROBLEMS 

1.  Amount  of   copper   sulphate  that  can  be   formed  from 
5  grams  of  cuprite  ?     How  much  zinc  would  be  required  to 
precipitate  this  amount  of  copper  from  solution  ? 

2.  From  8.4  grams  of  argentite  how  much  silver  could  be 
obtained?     How  much  silver  chloride,  nitrate,  and  sulphate 
can  be  formed  from  the  above  amount  of  silver  ? 

3.  Weight  of  gold  which  combines  with  4  litres  of  chlorine 
to  form  gold  trichloride  ?     Weight  of  auric  chloride  formed 
from  the  above  amount  of  gold  ? 


CHAPTER   XXV 

LEAD,   TIN 
LEAD  (At.  Wt.  =  206.9) 

THERE  remain  two  elements  in  group  IV  which  have  not 
been  studied.  These  are  tin  and  lead.  Although  the  atomic 
weight  of  tin  is  less  than  that  of  lead,  its  chemistry  is  more 
complex,  and  it  will  be  considered  after  lead. 

Occurrence,  Preparation,  and  Properties  of  Lead.  —  Lead 
occurs  in  nature  in  a  number  of  compounds.  The  most  im- 
portant is  the  sulphide  PbS,  or  galena. 

Galena,  being  the  principal  ore  of  lead,  is  the  one  from  which 
most  of  the  lead  of  commerce  is  prepared.  Several  methods 
are  employed  to  obtain  the  lead  from  the  sulphide.  One 
method  is  to  roast  the  sulphide,  converting  it  into  the  oxide, 
and  then  reduce  the  oxide  with  carbon  :  — 


2  PbS  +  3  02  =  2  PbO  +  2  S02  ; 
PbO  +  C  =  CO  +  Pb. 


A  second  method  is  to  roast  the  sulphide  until  a  part  of  it  is 
converted  into  the  oxide,  and  then  heat  the  oxide  with  the 
sulphide  :  — 

PbS  +  2  PbO  =  3  Pb  +  S02. 

The  action  of  water  upon  lead  is  of  hygienic  importance. 
Pure  water  dissolves  lead  much  more  readily  than  ordinary 
impure  water,  such  as  that  in  springs,  rivers,  etc.  The  im- 
purities react  with  the  lead  and  form  a  coating  of  carbonate, 
sulphate,  etc.,  which  protects  the  metal  from  the  further  action 
of  the  water.  If  the  water  contains  free  carbon  dioxide  or 
organic  acids,  it  acts  upon  the  lead,  converting  it  into  the  salt 
of  the  acid  in  question.  The  hygienic  question  is,  whether  drink- 

307 


308  ELEMENTS   OF   INORGANIC   CHEMISTRY 

ing-water  should  be  conducted  through  lead  pipes.  All  things 
considered,  it  is  much  safer  not  to  use  them,  since  pipes  made 
of  certain  other  metals  are  practically  unacted  upon  by  water. 
Lead  is  readily  precipitated  from  its  salts  by  a  number  of  metals. 
This  is  especially  the  case  with  zinc  and  iron.  When  a  bar  of 
zinc  is  suspended  in  a  solution  of  a  lead  salt,  the  lead  is  thrown 
out  and  the  zinc  dissolves.  What  takes  place  is  a  transfer  of 
the  electrical  charge  from  the  lead  ion  to  the  zinc  atom,  con- 
verting the  former  into  an  atom  and  the  latter  into  an  ion:  — 

Pb,  N03,  N03  -f-  Zn  =  Zn,  N03,  N03  +  Pb. 

Oxides  of  Lead.  —  Lead  forms  a  number  of  compounds  with 
oxygen.  Lead  suboxide,  Pb20 :  Lead  oxide,  PbO,  is  also  known 
as  litharge  or  massicot. 

Minium,  or  red  lead,  Pb304,  is  formed  by  gently  heating  lead 
oxide  on  the  air  (to  300°-400°). 

Lead  sesquioxide  has  the  composition,  Pb203. 

Lead  dioxide,  Pb02,  is  formed  by  oxidizing  the  lower  oxides  of 
lead  by  the  action  of  nitric  acid  upon  minium. 

Chloride   of  Lead.  —  Lead  generally  forms  the  bivalent  ion 

Pb,  which  readily  combines  with  the  anions  of  acids,  forming 
salts  that  are  beautifully  crystallized. 

Lead  chloride,  PbCl2,  is  readily  formed  by  bringing  together 

lead  ions,  Pb,  and  chlorine  ions,  Cl :  — 

Pb,  N03,  N03  +  Cl,  Na  +  Cl,  Na  =  Na,  N03  +  Na,  N~03  +PbCl2. 

Lead  chloride  is  a  white,  crystalline  substance,  somewhat 
soluble  in  hot  water,  but  only  slightly  soluble  in  cold  water. 

Lead  Sulphide,  PbS,  occurs  in  nature,  as  we  have  already 
seen,  as  galena.  It  is  formed  whenever  bivalent  lead  ions 
come  in  contact  with  sulphur  ions  —  whenever  a  soluble  lead 
salt  is  treated  with  hydrogen  sulphide :  — 

Pb(N03)2  +  H2S  =  2  HN03  +  Pbsf 

Lead  Carbonate,  PbC03.  —  The  carbonate  of  lead  is  an  im- 
portant compound,  and  especially  the  basic  carbonates.  The 


LEAD,   TIN  309 

normal  carbonate  occurs  in  nature  as  cerussite,  and  is  isonior- 
phous  with  aragonite,  a  form  of  calcium  carbonate.  The 
normal  carbonate  is  formed  when  ammonium  carbonate  is  added 
to  a  solution  of  lead  nitrate  :  — 

Pb(N03)2  +  (NH4)2C03  =  2  NH4N03  +  PbC03. 

If  any  other  alkaline  carbonate  is  used,  such  as  sodium  car- 
bonate, a  basic  lead  carbonate  is  precipitated,  and  this  is 
extensively  used  as  a  pigment  under  the  name  of  white 
lead. 

The  old  Dutch  method  of  making  white  lead  consists  in 
placing  sheet  lead,  rolled  into  spirals,  in  porcelain  pots  con- 
taining a  little  vinegar.  The  latter  did  not  touch  the  lead. 
The  vessels  were  then  placed  in  horse-manure,  which  decom- 
posed and  furnished  the  necessary  amount  of  carbon  dioxide. 
The  lead  plates  became  covered  in  time  with  a  layer  of  basic 
lead  carbonate,  which  was  removed  mechanically. 

This  method,  which  consumed  much  time,  has  now  given 
place  to  some  extent  to  quicker  processes.  Normal  lead 
acetate  is  shaken  with  litharge  and  water,  when  the  basic 
acetate  is  formed.  This  is  then  treated  with  carbon  dioxide, 
when  basic  lead  carbonate  is  formed.  The  ordinary  white 
lead  which  is  used  as  a  pigment  is  a  mixture  of  basic  car- 
bonates of  different  composition. 

Lead  Acetate,  Pb(CH3COO)2 .  3  H20,  is  an  important  soluble 
salt  of  lead.  It  is  formed  by  the  action  of  acetic  acid  on  lead 
oxide  or  finely  divided  lead.  On  account  of  its  sweet  taste  it 
is  known  as  sugar  of  lead. 

TIN  (At.  Wt.  =  119.0) 

An  element  chemically  allied  to  lead,  but  differing  from  it 
in  many  respects,  is  tin.  This  metal  is  useful  on  account 
of  its  chemical  inactivity,  is  valuable  on  account  of  its 
properties,  and  expensive  because  it  does  not  occur  in  large 
quantities.  Tin  occurs  in  the  uncombined  condition  along 


310  ELEMENTS   OF  INORGANIC   CHEMISTRY 

with  gold,  but  chiefly  as  tin  dioxide,  Sn02,  known  as  tinstone 
or  cassiterite. 

Preparation  and  Properties  of  Tin.  —  The  sulphur,  arsenic, 
and  similar  impurities  are  removed  from  the  tin  ore  by  roast- 
ing, and  the  oxide  is  then  heated  with  carbon.  The  oxide  is 
readily  reduced  to  the  metal,  and  this  is  purified  by  repeated 
melting,  the  molten  tin  being  poured  off  from  the  less  easily 
fusible  alloys. 

Tin  is  light  in  color  and  quite  soft.  It  can  be  readily  ham- 
mered or  rolled  into  thin  sheets  known  as  tin-foil,  which  is 
used  for  covering  objects  to  protect  them  from  the  action  of 
air,  moisture,  etc.  It  is  crystalline,  and  the  movement  of  the 
crystals  over  one  another  produces  a  crackling  noise  known  as 
the  cry  of  tin. 

Alloys  of  Tin.  —  Molten  tin  dissolves  readily  in  most  of  the 
other  metals  in  the  molten  condition,  and  forms  a  large  num- 
ber of  alloys  with  them.  Some  of  these  are  very  important 
substances.  Soft  solder  is  an  alloy  of  tin  and  lead.  The  bronzes, 
as  has  already  been  stated,  are  alloys  of  tin  and  copper. 

The  Tin  Ions.  —  The  element  tin  'forms  two  kinds  of  ions, 

+-)- 
those  which  are  bivalent,  Sn,  and  those  which  are  quadriva- 

++++ 
lent,  Sn.     Of  these  the  quadrivalent  ion  is  the  more  stable,  the 

bivalent  tending  to  pass  over  into  it.  No  other  ion  of  tin  is 
known.  The  stannous  ions  passing  over  into  the  stannic,  readily 
take  up  two  electrical  charges,  or  are  good  reducing  agents, 
as  we  say. 

Stannous  (SnO),  and  Stannic  (Sn02),  Oxides. —  The  two 
oxides  corresponding  to  the  stannous  and  stannic  conditions 
are  known.  Stannous  oxide,  SnO,  formed  by  heating  stannous 
hydroxide  in  a  current  of  carbon  dioxide,  is  comparatively 
unstable,  readily  combining  with  oxygen  and  forming  stannic 
oxide,  Sn02.  Stannic  oxide,  Sn02,  occurs  in  nature  as  tin-stone. 

Stannous  Hydroxides,  Sn(OH)2.  —  Stannous  hydroxide  is  pre- 
cipitated from  stannous  salts  by  the  addition  of  an  alkali :  — 

SnCl2  +  2  NaOH  =  2  NaCl  +  Sn(OH)2. 


LEAD,    TIN  311 

Stannous  hydroxide  readily  dissolves  in  an  excess  of  the  alkali, 
forming  salts. 

Stannic  Hydroxide,  Sn(OH)4,  is  formed  by  treating  tin  with 
nitric  acid  of  medium  concentration.  It  readily  loses  water, 
forming  metastannic  acid,  H2Sn03. 

Stannous  Chloride,  SnCl2  .  2  H20.  —  The  best-known  stannous 
salt  is  the  chloride,  SnCl2.  It  is  formed  by  dissolving  tin  in 
hydrochloric  acid,  also  by  heating  tin  in  a  current  of  dry 
hydrochloric  acid  gas.  It  combines  readily  with  chlorine, 
forming  stannic  chloride.  So  great  is  this  power  that  it  re- 
moves the  chlorine  from  the  chlorides  of  mercury,  precipitating 
metallic  mercury  :  — 

2  HgCl  +  SnClo  =  2  Hg  +  SnCl4  ; 
HgCl2  +  SnCl2  =  Hg  +  SnCl4. 


Stannous    chloride   combines    directly   with   free   chlorine, 
forming  stannic  chloride:  — 


Sn,  Cl,  Cl  +  C12  =  Sn,  01,  01,  01,  01. 

Stannous  chloride  is  known  commercially  as  tin-salt. 

Stannic  Chloride,  SnCl4.  —  The  tetrachloride  is  formed,  as  we 
have  seen,  by  the  action  of  mercurous  or  mercuric  chloride 
on  stannous  chloride.  It  is  also  formed  by  the  direct  action 
of  chlorine  gas  on  tin.  When  brought  in  contact  with  a  little 
water,  it  forms  a  viscous  mass  known  as  tin-butter,  having 
the  composition  SnCl4  .  3  H20. 

Sulphides  of  Tin.  —  Stannous  sulphide,  SnS,  is  formed  by 
conducting  hydrogen  sulphide  into  a  solution  of  stannous 
chloride  :  — 

SnCl2  +  H2S  =  2  HOI  +  SnS. 

Stannous  sulphide  is  a  brownish  black  powder,  insoluble  in 
dilute  acids,  but  soluble  in  ammonium  poly  sulphide,  forming 
ammonium  sulphostannate,  in  which  the  tin  is  in  the  quad- 
rivalent condition  :  — 

SnS  +  (NH4)2S2  =  (NH4)2SnS3. 


312  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Stannic  sulphide  is  precipitated  as  a  light  yellow  powder 
when  hydrogen  sulphide  is  passed  into  a  solution  of  a  stannic 
salt  or  a  stannate  :  — 

SnCl4  +  2  H2S  =  4  HC1  +  SnS2. 

Stannic  sulphide  dissolves  readily  in  ammonium  sulphide, 
forming  sulphostannates  or  thiostannates :  — 

SnS2  +  (NH4)2S  =  (NH4)2SnS3. 

This  reaction  is  of  importance  in  the  separation  of  tin  from 
most  of  the  elements. 


EXPERIMENTS  WITH  LEAD  AND  TIN 

Experiment  193.  Precipitation  of  Lead  from  its  Salts  by 
Other  Metals.  —  (Lead  nitrate  or  acetate ;  bar  of  zinc.) 

Add  a  few  grams  of  lead  nitrate  or  acetate  to  a  flask  holding 
2  litres  of  water  and  fill  the  flask  with  water,  adding  a  drop  or 
two  of  nitric  acid.  Suspend  a  bar  of  zinc  in  the  solution  of  the 
lead  salt.  The  lead  will  be  gradually  thrown  out  arid  deposited 
upon  the  zinc  in  arborescent  forms,  known  as  the  lead  tree. 

Allow  this  to  stand  several  days,  and  often  beautiful  results 
are  secured.  The  zinc  takes  the  electrical  charges  from  the 
lead  ions,  converting  them  into  atoms ;  while  the  zinc  atoms, 
having  received  electrical  charges,  become  ions  and  pass  into 
solution.  (See  page  308.) 

Experiment  194.  Oxides  of  Lead.  —  (Litharge ;  minium ;  lead 
sesquioxide ;  lead  dioxide.) 

Compare  the  several  oxides  of  lead  with  respect  to  their 
visible  properties.  Treat  minium  with  dilute  nitric  acid.  Is 
there  any  change  in  the  appearance  of  the  compound  ?  Is  it 
possible  to  dissolve  minium  completely  in  nitric  acid  ?  Filter 
the  solution,  preserving  both  the  residue  and  filtrate.  Test  the 
filtrate  for  lead  by  passing  in  hydrogen  sulphide  ? 

Treat  lead  dioxide  with  a  little  hydrochloric  acid  and  boil. 
What  is  given  off  ? 

Treat  lead  dioxide  with  hydrogen  dioxide  in  the  presence  of 
a  little  nitric  acid.  Explain  what  takes  place  ? 

Perform  these  same  experiments  with  the  residue  from 
minium  ?  (See  page  308.) 


LEAD,   TIN  313 

Experiment  195.  Lead  Chloride.  —  (Lead  nitrate;  sodium 
chloride.) 

Treat  a  solution  of  lead  nitrate  with  a  solution  of  sodium 
chloride.  White  lead  chloride  is  precipitated.  Boil  the  solu- 
tion, and  filter  hot.  Test  a  little  of  the  filtrate  for  lead  by  pass- 
ing in  hydrogen  sulphide.  Set  the  solution  aside  and  allow  it 
to  cool.  Lead  chloride  will  separate  in  the  form  of  crystals. 
Lead  chloride  is  only  slightly  soluble  in  cold  water,  but  much 
more  soluble  in  hot  water.  (See  page  308.) 

Experiment  196.  Lead  Sulphide.  —  (Lead  nitrate ;  hydrogen 
sulphide;  ammonium  sulphide.) 

Add  ammonium  sulphide  to  a  solution  of  lead  nitrate.  What 
is  precipitated  ?  Pass  hydrogen  sulphide  through  a  solution 
of  lead  nitrate.  Whatsis  precipitated?  Is  lead  sulphide  soluble 
in  dilute  acids  ?  (See  page  308.) 

Experiment  197.  Tin  thrown  down  by  Other  Metals. — 
(Stannous  chloride  ;  bar  of  zinc.) 

Suspend  a  bar  of  zinc  in  a  solution  of  stannous  chloride  to 
which  a  few  drops  of  hydrochloric  acid  have  been  added.  Allow 
the  zinc  to  stand  in  the  solution  for  some  days  in  a  quiet  place. 
What  is  formed  ? 

Zinc  takes  the  charge  from  the  tin  ions,  converting  them  into 
atoms.  The  zinc,  having  become  charged,  passes  into  solutions 
as  ions :  — 

Sn,  01,  Cl  +  Zn  =  Zn,  01,  01"  +  Sn. 

Experiment  198.  Hydroxides  of  Tin.  —  (Stannous  chloride ; 
sodium  hydroxide  ;  tin ;  concentrated  nitric  acid.) 

Add  sodium  hydroxide  to  a  solution  of  stannous  chloride. 
Stannous  hydroxide  is  thrown  down.  Equation?  Add  an 
excess  of  sodium  hydroxide.  Does  stannous  hydroxide  dissolve 
in  an  excess  of  sodium  hydroxide  ?  What  is  formed  ? 

Treat  a  few  grams  of  tin  with  a  little  concentrated  nitric 
acid.  Dense  fumes  are  given  off,  and  stannic  hydroxide  minus 
water,  H2Sn03,  is  formed :  — 

Sn(OH)4  -  H20  =  H2Sn03. 

This  compound,  which  is  a  white,  insoluble  powder,  is  known 
as  metastannic  acid.  (See  page  311.) 

Experiment  199.  Preparation  and  Properties  of  Stannous 
Chloride.  —  (Tin  ;  hydrochloric  acid.) 


314  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Treat  a  piece  of  tin  with,  hydrochloric  acid  and  warm.  Tin 
dissolves,  and  stannous  chloride  is  formed. 

Treat  a  little  of  the  stannous  chloride  solution  with  a  few 
drops  of  a  solution  of  mercuric  chloride.  The  mercuric  chloride 
is  reduced  to  white,  mercurous  chloride,  which,  being  insoluble, 
is  precipitated.  Add  more  of  the  solution  of  stannous  chloride 
and  warm  the  solution.  The  reduction  goes  farther ;  the  mer- 
curous chloride  being  reduced  to  metallic  mercury,  which  is 
thrown  down  as  a  black  precipitate.  (See  page  311.) 

Experiment  200.  Sulphides  of  Tin.  —  (Stannous  chloride ; 
stannic  chloride;  freshly  prepared  ammonium  sulphide ;  yellow 
sulphide  of  ammonium.) 

Pass  hydrogen  sulphide  into  a  solution  of  a  stannous  salt; 
also  into  a  solution  of  a  stannic  salt.  Stannous  sulphide  is 
thrown  down  in  the  former  case  and  stannic  sulphide  in  the 
latter. 

Treat  stannous  sulphide  with  freshly  prepared  ammonium 
sulphide.  Does  it  dissolve  ?  Treat  it  with  yellow  sulphide  of 
ammonium.  Does  it  dissolve  ? 

Treat  stannic  sulphide  with  freshly  prepared  ammonium 
sulphide.  Does  it  dissolve  ?  Treat  it  with  yellow  sulphide  of 
ammonium.  Does  it  dissolve  ?  How  would  you  distinguish  a 
stannous  from  a  stannic  compound  ?  (See  page  311.) 


PROBLEMS 

1.  Zinc  required  to  precipitate  206.9  grams  of  lead  from  a 
solution  of  lead  nitrate?     What  relation  does  this  bear  to  the 
atomic  weight  of  zinc? 

2.  Percentage  of  lead  in  each  of  the  four  oxides  of  lead  ? 

3.  Amounts  of  stannous  and  stannic  chlorides  that  can  be 
prepared  from  1  gram  of  stannous  oxide?  from  1  gram  of  stannic 
oxide  ? 


CHAPTER   XXVI 

I.    RUTHENIUM         RHODIUM         PALLADIUM 
II.    OSMIUM  IRIDIUM  PLATINUM 

ALTHOUGH  classed  together,  these  two  groups  of  three  ele- 
ments each  have  certain  properties  which  are  markedly  differ- 
ent. The  most  striking  is  the  atomic  weights.  The  first  three 
have  atomic  weights  which  are  not  widely  removed  from  one 
hundred,  while  the  atomic  weights  of  the  last  three  are  not 
widely  removed  from  two  hundred.  Similarly,  the  specific 
gravities  of  the  first  three  elements  is  close  to  twelve,  while  the 
specific  gravities  of  the  last  three  is  about  twenty-two. 

These  elements,  however,  all  resemble  platinum  more  or  less 
closely,  and  are  all  comparatively  rare;  platinum  itself  not 
occurring  in  very  large  quantities. 

The  chemistry  of  these  substances  presents  certain  points  of 
interest.  Palladium  absorbs  large  volumes  of  hydrogen  gas, 
and  it  has  been  supposed  for  a  long  time,  forms  the  compound 
Pd2H,  known  as  palladium  hydride.  There  is  some  doubt  at 
present  as  to  whether  this  is  a  definite  chemical  compound. 

Osmium  combines  with  oxygen,  forming  the  compound  Os04. 
This  compound  is  remarkable  in  that  there  is  no  element 
known  which  has  the  power  to  combine  with  more  than  four 
oxygen  atoms. 

PLATINUM   (At.  Wt.  =  195.0) 

Platinum,  one  of  the  most  valuable  elements  from  the  stand- 
point of  chemistry,  on  account  of  its  power  to  resist  chemical 
reagents,  occurs  fairly  widely  distributed,  but  not  in  large 
quantities.  It  occurs  in  company  with  the  other  platinum 
metals,  and  is  separated  from  them  by  methods  with  which  we 
are  now  more  or  less  familiar. 

315 


316  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Properties  of  Platinum.  —  Platinum  is  light  in  color,  has  a 
specific  gravity  of  21.4,  and  melts  at  about  1770°  in  the  flame 
of  the  oxy hydrogen  blowpipe.  It  is  both  ductile  and  malleable 
—  can  be  drawn  into  thin  wire  and  hammered  into  thin  sheets. 

Platinum  can  exist  in  a  number  of  physical  conditions.  In 
addition  to  ordinary  white  platinum,  which  is  very  compact 
and  metallic  in  all  of  its  properties,  we  know  platinum  in  the 
form  of  a  spongy  mass  which  is  known  as  platinum  sponge. 
Platinum  is  obtained  in  this  condition  when  the  chloride  of 
ammonium  and  platinum,  which  will  be  referred  to  a  little 
later,  is  heated.  Platinum  sponge  has  a  gray  color,  and  differs 
fundamentally  from  the  black  variety  of  platinum  known  as 
platinum  black.  This  is  obtained  by  reduction  of  platinum 
compounds.  When  both  of  these  varieties  are  highly  heated 
and  subjected  to  pressure,  they  pass  back  into  ordinary  white 
platinum.  Finely  divided  platinum  has  a  remarkable  power 
to  absorb  oxygen.  It  can  absorb  several  hundred  volumes  of 
oxygen,  and  the  oxygen  in  this  condition  is  very  active  chemi- 
cally, readily  effecting  oxidations.  In  a  similar  manner  it  can 
absorb  considerable  quantities  of  hydrogen,  and  the  hydrogen 
under  these  conditions  has  strongly  reducing  properties. 

This  power  of  platinum  to  absorb  gases,  and  to  produce 
chemical  combinations  between  them,  can  be  readily  illustrated 
by  means  of  a  piece  of  platinum  foil  and  a  gas-jet.  Ignite  the 
jet  and  heat  the  platinum  to  re'dness.  Then  extinguish  the  jet 
until  the  platinum  ceases  to  glow.  If  now  the  gas-jet  is  turned 
on  again  and  the  gas  allowed  to  flow  over  the  surface  of  the 
hot,  but  not  incandescent,  platinum,  it  will  be  reignited.  This 
experiment  can  be  repeated  as  often  as  desired  with  the  same 
piece  of  platinum. 

The  power  of  platinum  to  absorb  gases  and  produce  chemi- 
cal reaction  catalytically,  is  also  illustrated  by  a  form  of  lamp 
devised  by  Davy. 

An  ordinary  spirit  lamp  is  filled  with  a  mixture  of  alcohol 
and  ether,  and  ignited.  A  spiral  of  platinum  wire  is  suspended 
in  the  flame  and  heated  to  incandescence.  The  flame  is  then 


PLATINUM  317 

extinguished  until  the  spiral  ceases  to  glow.  If  'the  vapors  of 
alcohol  and  ether  are  now  allowed  to  fall  again  on  the  hot 
platinum,  the  latter  will  again  become  incandescent  and  ignite 
the  lamp. 

The  Dobereiner  lamp,  constructed  early  in  the  nineteenth 
century,  is  based  upon  the  same  principle.  A  current  of  hy- 
drogen is  allowed  to  flow  over  spongy  platinum,  when  it  com- 
bines so  rapidly  with  oxygen  that  the  platinum  is  heated  to 
incandescence  and  ignites  the  hydrogen. 

Chemically,  platinum  is  very  resistant  to  reagents,  and  upon 
this  fact  its  value  chiefly  depends.  It  is  not  attacked  appre- 
ciably by  any  of  the  strong  mineral  acids,  but  dissolves  in  aqua 
regia.  It,  however,  dissolves  in  the  fused  alkalies,  cyanides, 
etc.  Platinum  must  not  be  heated  in  contact  with  carbon, 
since  some  of  the  carbon  dissolves  and  makes  the  platinum 
brittle. 

Uses  of  Platinum.  —  Platinum  is  especially  useful  to  the 
chemist  on  account  of  its  resistance  to  chemical  reagents.  It 
is  not  attacked  appreciably  even  by  hydrofluoric  acid.  It  is 
also  not  attacked  by  the  fused  alkaline  carbonates.  Platinum 
vessels  can  thus  be  used  where  even  porcelain  could  not  be 
employed.  On  account  of  its  high  melting-point  platinum  is 
used  for  holding  substances  which  are  to  be  heated  to  a  high 
temperature.  Since  it  does  not  volatilize  in  the  flame  of  the 
bunsen  burner,  it  does  not  impart  any  color  to  the  flame,  and 
is  therefore  useful  in  spectrum  analysis  to  hold  the  substance 
whose  spectrum  it  is  desired  to  study.  A  platinum  wire  is 
made  into  a  loop  and  dipped  into  the  substance  in  question. 
It  is  then  inserted  into  the  flame,  when  only  the  spectrum 
characteristic  of  the  substance  appears. 

Platinum  finds  to-day  extensive  applications  in  the  arts. 
Platinum  has  nearly  the  same  temperature  coefficient  of  ex- 
pansion as  glass,  and  almost  exactly  the  same  as  red  and 
blue  fusion  glass.  If  it  is  desired  to  seal  an  electrical  connec- 
tion through  glass,  a  platinum  wire  is  the  most  convenient 
means.  This  fact  is  utilized  not  only  in  constructing  standard 


318  ELEMENTS   OF   INORGANIC   CHEMISTRY 

cells,  but  also  incandescent  electric  lights.  Large  amounts  of 
platinum  are  used  in  this  way. 

Oxides  and  Hydroxides  of  Platinum.  —  Platinum  forms  two 
oxides,  PtO  and  Pt02,  which  are  derived  from  the  correspond- 
ing hydroxides,  Pt(OH)2  and  Pt(OH)4,  by  careful  heating. 
The  hydroxides  are  obtained  from  the  platinous  and  plantinic 
chlorides,  or  from  the  double  chlorides  with  the  alkalies,  by 
the  action  of  a  base.  All  of  these  substances  are  unstable 
at  elevated  temperatures,  breaking  down  and  yielding  metallic 
platinum. 

Chlorides  of  Platinum. —  Platinum  forms  two  chlorides, — 

++ 

platinous  chloride,  PtCl2,  in  which  the  platinum  is  bivalent,  Ft, 

++++ 
and  plantinic  chloride,  in  which  the  platinum  is  tetravalent,  Pt. 

Platinic  chloride,  which  is  quite  soluble  in  water,  readily  com- 
bines with  hydrochloric  acid,  forming  Jiydrochlorplatinic  acid, 
H2PtCl6.  This  same  compound  is  formed  when  platinum  is 
dissolved  in  aqua  regia. 

Many  salts  of  this  acid  are  known.  The  potassium  and 
ammonium  chlorplatinates,  K2PtCl6  and  (NH4)2PtCl6,  are 
difficultly  soluble  in  water,  while  the  sodium  salt,  Na2PtCl6,  is 
readily  soluble.  This  reagent  thus  enables  us  to  separate 
sodium  from  potassium. 

In  this  acid  the  chlorplatinic  ion,  PtCl6,  is  bivalent,  and 
this  is  an  example  of  a  metal  forming  part  of  an  anion.  When 
a  solution  of  this  acid  is  electrolyzed,  the  platinum  passes  to  the 
anode  and  not  to  the  cathode. 


CHAPTER    XXVII 

GENERAL    RELATIONS 

RELATIONS   WITHIN    THE    GROUPS   OF   THE    PERIODIC 

SYSTEM 

Group  I.  —  The  chemical  relations  between  sodium  and  po- 
tassium are  very  close  indeed.  These  elements  have  the  same 
general  properties  and  show  the  same  reactions  towards  water 
and  other  reagents.  The  hydroxides  are  both  univalent,  and 
are  the  strongest  bases.  The  salts  with  all  of  the  more  com- 
mon acids  are  surprisingly  similar,  and  these  elements  are  as 
closely  allied  as  any  two  of  the  more  common  elements  in  the 
whole  field  of  chemistry. 

Sodium  and  potassium  are  also  very  closely  allied  to  the 
rarer  elements,  lithium,  rubidium,  and  caesium. 

Copper,  silver,  and  gold,  while  bearing  many  relations  to  one 
another,  differ  markedly  and  in  many  respects  from  the  alkali 
metals  just  considered.  Copper  may  be  univalent  and  bivalent, 
and  gold  univalent  and  trivalent.  These  substances  are  not 
such  strong  base-forming  elements  as  the  alkalies.  Indeed,  the 
power  to  form  acids  begins  to  manifest  itself  in  gold. 

Group  II.  —  Calcium,  strontium,  and  barium  are  very  closely 
allied.  Indeed,  about  as  closely  as  sodium  and  potassium. 
These  elements  all  decompose  water,  forming  strong  bases. 
They  are  all  bivalent,  and  their  compounds  with  the  different 
acids  are  strikingly  alike.  These  compounds  may  differ  in 
their  solubility  in  certain  solvents,  but,  in  general,  when  one 
of  the  alkaline  earths  forms  a  difficultly  soluble  salt  with  an 
acid,  the  corresponding  salts  of  the  other  two  elements  are 
difficultly  soluble.  Magnesium,  zinc,  cadmium,  and  mercury 
are  more  closely  allied  to  each  other  than  to  calcium,  strontium, 

319 


320  ELEMENTS   OF   INORGANIC  CHEMISTRY 

and  barium.  Magnesium,  however,  resembles  the  calcium 
group  much  more  closely  than  the  others. 

Zinc  is  always  bivalent,  but  it's  hydroxide  is  both  basic  and 
acidic,  depending  upon  conditions.  Its  sulphide  is  insoluble  in 
water,  and  thus  differs  from  all  preceding  elements.  Cadmium 
is  nearly  always  bivalent,  but  can  act  as  a  univalent  element, 
while  mercury  is  both  univalent  and  bivalent. 

Group  III.  —  The  first  element,  boron,  is  a  strongly  acid- 
forming  element,  while  aluminium  in  its  hydroxide  is  both  base- 
forming  and  acid-forming.  These  elements  are  both  trivalent. 
The  remaining  elements  of  this  group  are  comparatively  rare 
substances. 

Group  IV.  —  Carbon  and  silicon  are  closely  allied  in  many 
respects.  They  are  both  quadrivalent  in  most  of  their  reactions, 
and  are  both  markedly  acid-forming  elements.  Both  form 
homologous  series  of  compounds,  and  their  chemistry  is,  there- 
fore, comparatively  complex. 

Tin  and  lead,  the  remaining  two  abundant  elements  in  this 
group,  differ  markedly  from  the  first  two,  but  show  certain  not 
very  close  relations  to  one  another.  Tin  is  bivalent  and  quad- 
rivalent, while  lead  is  bivalent.  The  hydroxides  of  both  of 
these  elements  are  both  basic  and  acidic,  depending  upon  con- 
ditions. They  are  both  well-characterized  metals,  while  carbon 
and  silicon  are  pronounced  metalloids. 

Group  V.  —  The  relationships  in  group  V  are  very  satisfac- 
tory, at  least  as  far  as  the  well-known  elements  nitrogen,  phos- 
phorus, arsenic,  and  antimony  are  concerned.  The  elements 
in  this  group  have  variable  valence,  but  are  generally  trivalent 
or  pentavalent.  They  are  all  acid-forming  elements,  but  most 
of  them  can  combine  with  strong  acids  and  form  salts.  The 
compounds  of  these  several  elements  with  hydrogen,  and  with 
hydrogen  and  oxygen,  are  strikingly  alike  in  their  properties. 
Their  compounds  with  chlorine  also  closely  resemble  each  other. 

Bismuth,  which  has  the  highest  atomic  weight  of  any  mem- 
ber of  the  series,  shows  basic  properties. 

Group  VI.  —  Here  the  familiar  elements  oxygen  and  sulphur. 


GENERAL   RELATIONS  321 

and  the  rarer  elements  selenium  and  tellurium  have  many 
points  of  resemblance.  Their  compounds  with  hydrogen  have 
strictly  analogous  composition.  They  are  all  strongly  acid- 
forming  elements,  and  all  except  oxygen  have  a  valence  as 
high  as  six. 

The  remaining  three  elements  in  this  group  are  comparatively 
rare  substances,  and  cannot  be  brought  within  the  scope  of 
this  work. 

Group  VII.  —  The  relationships  here  between  fluorine,  chlo- 
rine, bromine,  and  iodine  are  as  close  or  closer  than  in  any 
other  group.  Their  compounds  with  hydrogen,  with  hydrogen 
and  oxygen,  with  the  metals  and  non-metals  in  general,  show 
surprisingly  close  relations.  They  are  among  the  strongest 
acid-forming  elements,  and  have  a  maximum  valence  of  seven. 

The  relation  of  manganese  to  the  halogens  is  not  so  obvious. 
There  are  certain  analogies  between  the  compounds  of  man- 
ganese and  those  of  the  halogens.  Thus,  permanganic  acid 
resembles  in  composition  perchloric  acid,  the  manganese  having 
a  valence  of  seven.  If,  however,  we  consider  in  a  broad  way 
the  chemistry  of  manganese,  we  shall  find  it  much  more  closely 
allied  to  iron  and  chromium  than  to  the  halogens,  and  it  is 
a  weak  point  in  the  Periodic  System  that  it  should  fall  in 
group  VII. 

Group  VIII. — This  group  contains  those  elements  which 
are  usually  not  regarded  as  fitting  into  the  system.  Certain 
close  relations,  however,  exist  between  the  several  sub-groups 
of  three  elements  within  this  group. 

Iron,  cobalt,  and  nickel  show  many  relations  both  in  the 
elementary  condition  and  in  their  hydroxides  and  salts. 

The  rare  elements  ruthenium,  rhodium,  and  palladium  are 
also  closely  allied ;  while  osmium,  iridium,  and  platinum  are 
among  the  most  closely  allied  of  the  chemical  elements.  Even 
within  this  overflow  group  we  have,  then,  many  unquestionable^ 
relations  manifesting  themselves. 


322  ELEMENTS  OF  INORGANIC  CHEMISTRY 

RELATIONS  BETWEEN  THE   COMPOUNDS   OF   THE 
METALS 

Oxides.  — The  oxides  of  the  metals,  with  the  exception  of 
the  alkalies,  are  difficultly  soluble  in  water.  The  oxides  of 
the  alkaline  earths  are  somewhat  soluble  in  water.  Indeed, 
most  of  the  oxides  are  practically  insoluble,  and  occur  in 
nature  as  valuable  ores.  This  is  the  case  especially  with 
aluminium,  iron,  manganese,  and  the  like. 

Hydroxides.  —  The  hydroxides  of  the  metals,  with  the  ex- 
ception of  the  alkalies,  are  generally  difficultly  soluble  in 
water.  This  has  been  shown  by  repeated  experiments  in 
which  the  hydroxides  of  the  heavy  metals  were  formed  by 
treating  the  soluble  salts  of  the  metals  with  a  soluble  hydrox- 
ide. The  hydroxides  of  the  alkaline  earths  are  somewhat 
soluble  in  water  —  barium  being  the  most  soluble  and  calcium 
the  least. 

Certain  of  the  hydroxides,  like  zinc,  aluminium,  and  lead, 
are  of  special  interest  in  that  they  act  as  bases  towards  the 
strong  acids,  and  as  acids  towards  the  stronger  bases.  The 
basic  nature  of  the  hydroxides  is  most  pronounced  with  the 
alkalies,  and  becomes  less  and  less  as  we  pass  to  the  right  in 
the  Periodic  System,  as  has  already  been  pointed  out. 

Halides.  —  The  halides  of  the  metals  include  the  chlorides, 
bromides,  iodides,  and  fluorides.  They  are  formed  by  dissolv- 
ing the  metal,  the  oxide,  or  the  carbonate  in  the  acid  in  ques- 
tion. The  chlorides  are  formed  by  treating  soluble  sulphates 
with  barium  chloride.  Write  the  equations  in  each  case  ? 

The  chlorides  of  most  of  the  metals  are  soluble  in  water. 
Silver  and  mercurous  chlorides  are  exceptions.  Lead  chloride 
is  quite  soluble  in  hot  water,  but  soluble  with  difficulty  in  cold 
water.  These  facts  are  important,  as  we  shall  see,  in  connec- 
tion with  the  qualitative  separations  of  the  metals. 

Nitrates.  —  The  nitrates  of  the  metals  are  formed  by  dissolv- 
ing the  metal,  the  oxide,  or  the  carbonate  in  nitric  acid.  Also 
by  treating  soluble  sulphates  or  carbonates  with  the  nitrates  of 


GENERAL   RELATIONS  323 

elements  which  form  insoluble  sulphates  or  carbonates.  Thus, 
when  sodium  sulphate  is  treated  with  barium  nitrate,  barium 
sulphate  is  precipitated  and  sodium  nitrate  remains  in 
solution. 

The  nitrates  of  all  the  metals  are  soluble  in  water.  These 
easily  give  up  oxygen  at  elevated  temperatures  and  are,  there- 
fore, good  oxidizing  agents. 

Sulphates.  —  The  sulphates  are  formed  by  the  solution  of 
metals,  oxides  of  the  metals,  or  carbonates  of  the  metals  in 
sulphuric  acid.  When  the  sulphate  is  insoluble,  it  is  formed 
by  treating  a  soluble  salt  of  the  metal  with  a  soluble  sulphate, 
or  with  sulphuric  acid.  The  sulphates  of  the  alkalies  are 
fairly  soluble,  but  are  not  to  be  classed  among  the  very  soluble 
compounds.  The  sulphates  of  the  alkaline  earths  are  very 
insoluble  substances.  Calcium  sulphate,  to  be  sure,^  dissolves 
slightly  in  water,  but  strontium  sulphate  is  less  soluble,  and 
barium  sulphate  one  of  the  most  insoluble  compounds 
known  in  the  whole  field  of  chemistry. 

All  degrees  of  solubility  are  represented  among  the  sulphates 
of  the  heavy  metals.  In  general,  they  are  more  soluble  than 
the  sulphates^  of  the  alkaline  earths. 

Carbonates.  —  Since  carbonic  acid  is  so  weak  and  is  volatile, 
the  carbonates  cannot  be  prepared  by  dissolving  the  metal  or 
its  compounds  in  a  solution  of  carbonic  acid.  The  carbonates 
are  prepared  by  passing  carbon  dioxide  through  a  solution  of 
the  hydroxide  of  the  metal  in  question.  Also  by  the  action  of 
a  soluble  carbonate  on  a  solution  of  a  salt  of  the  metal. 

The  carbonates  of  the  alkalies  are  soluble.  The  carbonates 
of  the  other  elements  are  insoluble.  Some  of  the  carbonates 
occur  in  nature  in  large  quantities.  This  is  especially  true  of 
the  carbonates  of  calcium  and  magnesium.  The  acid  carbon- 
ates are  often  more  soluble  than  the  carbonates. 

The  carbonates  of  the  metals  are  decomposed  by  heat,  yield- 
ing carbon  dioxide  which  is  volatile,  and  the  oxide  of  the  metal. 
The  temperature  required  varies  greatly  from  one  carbonate  to 
another. 


324  ELEMENTS   OF  INORGANIC   CHEMISTRY 

Phosphates.  —  The  phosphates  are  formed  when  soluble  salts 
of  the  metals  are  treated  with  a  solution  of  a  soluble  phos- 
phate ;  e.g.  with  disodium  phosphate.  The  normal  phosphates 
of  all  of  the  metals,  with  the  exception  of  the  alkalies,  are 
insoluble  in  water.  The  secondary  and  primary  phosphates 
may  differ  greatly  from  the  normal  phosphate  in  solubility. 
Such  is  the  case  with  the  various  phosphates  of  calcium. 
Some  of  the  phosphates  occur  in  nature  as  minerals. 

Sulphides.  —  The  sulphides  are  specially  interesting  in  con- 
nection with  the  qualitative  separations  of  the  metals  from  one 
another.  There  are  some  metals  whose  sulphides  are  soluble 
in  water;  others  whose  sulphides  are  soluble  in  dilute 
acids;  others  again  whose  sulphides  are  insoluble  in  dilute 
acids ;  and  still  others  whose  sulphides  dissolve  in  ammonium 
polysulphide.  Upon  these  facts  is  based  the  division  of  the 
elements,  from  the  standpoint  of  qualitative  analysis,  into 
several  groups. 

Group  I  includes  those  elements  whose  chlorides  are  in- 
soluble in  water,  and  which  are,  therefore,  precipitated  from 
solutions  of  their  soluble  salts  by  soluble  chlorides  or  hydro- 
chloric acid.  These  include  silver,  lead,  and  mercurous 
mercury. 

Leaving  group  I,  or  those  metals  whose  chlorides  are  insoluble 
in  water  (lead,  silver,  and  mercurous  mercury),  we  come  to 
Group  II,  or  those  elements  whose  sulphides  are  precipitated  by 
hydrogen  sulphide  in  the  presence  of  dilute  acid,  and  whose 
sulphides  are  soluble  in  the  polysulphides  of  ammonium. 
These  are  arsenic,  antimony,  and  tin,  and  also  gold  and 
platinum. 

Group  III  includes  those  elements  whose  sulphides  are  pre- 
cipitated by  hydrogen  sulphide  in  the  presence  of  a  dilute 
acid,  and  whose  sulphides  are  insoluble  in  ammonium  poly- 
sulphides.  These  include  cadmium,  mercury,  copper,  lead, 
silver,  and  bismuth. 

Group  IV  includes  those  metals  whose  sulphides  are  not  pre- 
cipitated by  hydrogen  sulphide,  since  their  sulphides  are 


GENERAL  RELATIONS  325 

soluble  in  dilute  acids,  but  are  precipitated  by  ammonium 
sulphide  either  as  sulphides  or  hydroxides.  The  members  of 
this  group  are  iron,  cobalt,  nickel,  zinc,  and  manganese,  which 
are  precipitated  as  the  sulphides ;  and  aluminium  and  chromium, 
which  are  precipitated  as  the  hydroxides. 

Group  V  contains  those  elements  whose  sulphides  are  not 
precipitated  by  ammonium  sulphide,  but  whose  solutions  are 
precipitated  by  a  soluble  carbonate.  These  are  calcium, 
barium,  strontium,  and  magnesium. 

Group  VI  contains  the  alkalies:  sodium,  potassium,  and 
ammonium.  On  account  of  the  solubility  of  their  salts  they 
are  not  precipitated  from  their  solutions  by  any  of  the  ordi- 
nary reagents.  To  precipitate  them,  such  reagents  as  platinic 
chloride,  perchloric  acid,  tartaric  acid,  etc.,  are  required. 

Separation  of  the  Members  of  Any  Group.  —  To  separate  the 
individual  members  of  any  group  from  one  another,  special 
reactions  characteristic  of  the  elements  in  question  must  be 
used.  It  would  lead  too  far  to  give  a  complete  scheme  of 
qualitative  analysis  in  this  place,  and  any  scheme  which  is  not 
complete  is  not  satisfactory.  In  this  connection  reference 
must  be  had  to  some  one  of  the  many  excellent  works  on  quali- 
tative analysis. 

THE  NATURE  AND  ROLE   OF  IONS  IN  CHEMISTRY 

We  can  see  from  what  we  have  learned  that  the  metals  in 
general  form  cations,  while  the  acid  radicles  form  anions.  A 
chemical  compound  is  formed  by  the  union  of  an  anion  with 
a  cation.  In  some  cases  the  metal,  instead  of  forming  a  cation, 
unites  with  other  elements  forming  part  of  the  anion,  as  plati- 
num in  potassium  platinum  chloride.  It  is,  however,  a  very 
simple  matter  in  general  to  determine  at  a  glance  which  con- 
stituent of  a  compound  is  the  cation  and  which  the  anion.  The 
most  positive  constituent,  or  strongest  base-forming  element, 
as  it  is  said,  is  the  cation,  and  the  remainder  of  the  compound 
the  anion. 


326  ELEMENTS   OF  INORGANIC   CHEMISTRY 

We  thus  see  that  the  cations  of  compounds  are  in  general 
very  simple,  being  in  most  cases  atoms  of  the  metals ;  while 
the  anion  may  be  simple,  as  the  anion  of  hydrochloric  acid 
Cl,  or  may  be,  and  usually  is,  complex,  as  the  anion  of  sulphuric 

acid,  S04. 

The  importance  of  the  ions  in  chemistry  has  come  to  be 
recognized  in  the  last  few  years.  It  is  now  quite  certain  that 
most  reactions  in  inorganic  chemistry  are  reactions  between 
ions,  molecules  as  such  having  nothing  to  do  with  the  reac- 
tion. The  molecules  dissociate,  yielding  the  ions,  which  then 
enter  into  chemical  reaction.  Some  little  of  the  evidence 
leading  to  this  conclusion,  which  is  of  fundamental  importance 
for  the  whole  science  of  chemistry,  has  already  been  referred 
to.  For  a  fuller  discussion  of  this  and  similar  subjects,  larger 
works  must  be  consulted. 


INDEX 


Accumulators,  161. 

Acetylene,  193,  242. 

Acetylene,  preparation  of ,  experiment, 
211. 

Acid  and  basic  properties  of  the  ele- 
ments, 131. 

Acid,  definition  of,  66. 

Acidity  of  bases  and  basicity  of  acids, 
99. 

Acids  and  bases,  experiments  with, 
104. 

Agate,  205. 

Air,  107. 

Air,  a  mixture  or  compound,  108. 

Air,  liquid,  109. 

Air,  physical  properties  of,  109. 

Air-slaked  lime,  241. 

Albite,  270. 

Alcohol  formed  by  the  fermentation 
of  sugar,  experiment,  214. 

Alcohols,  197. 

Alkalies,  are  base-forming  and  univa- 
lent  ions,  235. 

Alkali  metals,  220. 

Alkali  metals,  characteristics  of,  234. 

Alloys  of  gold,  304. 

Aluminate  of  sodium,  268. 

Aluminium  carbide,  269. 

Aluminium  chloride,  268. 

Aluminium  hydroxide  precipitated  by 
a  soluble  carbonate  or  by  ammo- 
nium sulphide,  experiment,  272. 

Aluminium  hydroxide  soluble  in  so- 
dium hydroxide,  experiment,  271. 

Aluminium,  occurrence  and  prepara- 
tion, 266. 

Aluminium  oxide  and  hydroxide,  267. 

Aluminium,  properties  of,  267. 

Aluminium  silicates,  applications  of, 
270. 


Aluminium,  silicates  of,  270. 

Aluminium  sulphate,  269. 

Alums,  269. 

Amalgamation  process,  299. 

Amalgams,  261. 

Amalgams,  formation  of,  experiment, 
264. 

Amethyst,  205. 

Amethyst,  oriental,  267. 

Ammonia,  absorption  by  charcoal, 
experiment,  91. 

Ammonia,  absorption  by  water,  ex- 
periments, 90. 

Ammonia,  chemical  properties  of,  78. 

Ammonia  dissolved  hi  water  is  a  base, 
experiment,  88. 

Ammonia,  gaseous,  reacts  with  gaseous 
hydrochloric  acid,  experiment,  89. 

Ammonia  in  the  atmosphere,  108. 

Ammonia  lighter  than  air,  experiment, 
90. 

Ammonia,  neutralization  of  by  acids, 
experiment,  89. 

Ammonia  or  Solvay  process  for  prepar- 
ing sodium  carbonate,  224. 

Ammonia,  physical  properties  of,  78. 

Ammonia,  preparation  from  ammo- 
nium chloride  and  slaked  lime, 
experiment,  87. 

Ammonia,  preparation  from  ammo- 
nium chloride  and  sodium  hydrox- 
ide, experiment,  87. 

Ammonia  when  neutralized  forms  a 
salt,  experiment,  89. 

Ammonium,  232. 

Ammonium  acid  carbonate,  234. 

Ammonium  amalgam,  79. 

Ammonium  and  sodium  acid  phos- 
phate, 234. 

Ammonium  carbonate,  234. 


327 


328 


INDEX 


Ammonium  chloride,  233. 

Ammonium  hydrosulphide,  233. 

Ammonium  hydroxide,  232. 

Ammonium  polysulphides,  233. 

Ammonium  salts,  volatility  of,  experi- 
ment, 237. 

Ammonium  sulphate,  234. 

Ammonium  sulphide,  233. 

Ammonium  sulphide  and  hydrosul- 
phide, formation  of,  experiment, 
237. 

Ammonium  sulphide  group,  153. 

Anions,  64. 

Annealing  glass,  247. 

Anode,  266. 

Anode  plates,  296. 

Anthracite,  191. 

Antimonic  acid,  180. 

Antimonious-sulph,  and  antimonic- 
sulph,  acids,  181. 

Antimony,  compounds  of  with  sul- 
phur and  the  metals,  180. 

Antimony,  fusion  of  combustion  of 
in  contact  with  the  air,  experiment, 
187. 

"  Antimony  mirror,"  187. 

Antimony,  occurrence  and  prepara- 
tion, 179. 

Antimony,  oxides  of,  180. 

Antimony  sulphide,  experiment,  187. 

Apatite,  172,  245. 

Aqua  regia,  85. 

Aragonite,  243. 

Argentan,  297. 

Argentite,  298. 

Argon,  110. 

Arsenic  acid,  179. 

Arsenical  pyrites,  178. 

Arsenic,  compounds  with  oxygen,  178. 

"Arsenic  mirror,"  186. 

Arsenic,  occurrence,  preparation,  and 
properties,  178. 

Arsenic  pentoxide,  179. 

Arsenic  sublimes  without  melting, 
burns  on  the  air,  formed  by  reduc- 
ing the  oxide  with  carbon,  experi- 
ment, 185. 

Arsenic,  sulpho-salts  of,  179. 

Arsenic  trioxide,  179. 

Arsenic  trisulphide,  experiment,  186. 

Arsenious  acid,  179, 


Arsine,  178. 

Arsine  decomposed  by  heat,  experi- 
ment, 186. 

Arsine  preparation  of,  experiment, 
185. 

Asbestos,  256. 

Atmosphere,  composition  of  the,  107. 

Atmospheric  air,  and  certain  rare  ele- 
ments occurring  in  it,  107. 

Atmospheric  air,  experiments  with, 
111. 

Atmospheric  air,  physical  properties 
of,  109. 

Atomic  weights  and  chemical  proper- 
ties, combining  power,  129. 

Atomic  weights  and  combining  num- 
bers, 115. 

Atomic  weights  and  properties  of  the 
elements,  relations  between,  126. 

Atomic  weights,  determination  of  rela- 
tive, 115. 

Atomic  weights  from  molecular 
weights,  119. 

Atomic  weights,  table  of,  121. 

Atoms,  relative  weights  of,  121. 

Available  phosphoric  acid,  246. 

Avogadro's  hypothesis,  116. 

Avogadro's  hypothesis  and  molecular 
weights,  117. 

Azurite,  296. 

Barite,  248. 

Barium,  248. 

Barium  dioxide,  249. 

Barium,  flame  reaction  of,  experi- 
ments, 254. 

Barium  hydroxide,  249. 

Barium  hydroxide,  properties  of,  ex- 
periment, 254. 

Barium  oxide,  249. 

Barium  sulphate,  249. 

Barium  sulphate  decomposed  by  so- 
dium carbonate,  experiment,  254. 

Basic  and  acid  properties  of  the  ele- 
ments, 131. 

Basicity  of  acids  and  acidity  of  bases, 
99. 

Basic  lining,  277. 

Bauxite,  267. 

Bell  metal,  297. 

Bessemer  converter,  276. 


INDEX 


329 


Bessemer  steel,  276. 

Bicarbonate  of  soda,  225. 

Bismuth,  compounds  with  oxygen  and 
hydrogen,  182. 

Bismuth,  occurrence  and  properties, 
181. 

Bismuth,  oxidation  of  in  the  arts,  ex- 
periment, 187. 

Bismuth  sulphide,  188. 

Bituminous  coal,  191. 

Bivalent  ion,  72. 

Blast-furnace,  273. 

Bleaching  agent,  45. 

Bleaching-powder,  69,  241. 

Bleaching-powder,  preparation  of, 
experiment,  74. 

Blowing,  glass,  247. 

Blowpipe,  202. 

Blowpipe,  oxyhydrogen,  22. 

Blowpipe,  use  of  the,  experiment, 
216. 

Blue  vitriol,  298. 

Bog-iron  ore,  273. 

Bone-ash,  245. 

Bone-black,  190. 

Borax,  208. 

Boric  acid,  208. 

Boric  acid,  preparation  of,  color  im- 
parted to  the  alcohol  flame,  experi- 
ment, 217. 

Boron,  207. 

Boron  nitride,  209. 

Boron,  occurrence,  preparation,  and 
properties,  208. 

Boron  trioxide,  208. 

Boyle's  law  for  gases,  112. 

Brass,  258, 297. 

Braunite,  285. 

Brick,  270. 

Brimstone,  crude,  150. 

Bromic  acid,  136. 

Bromine,  133. 

Bromine,  action  of  on  other  elements, 
experiment,  143. 

Bromine  atoms  and  bromine  ions, 
134. 

Bromine,  chemical  properties  of, 
134. 

Bromine,  compounds  with  oxygen 
and  hydrogen,  136. 

Bromine,  detection  of,  134. 


Bromine,  detection  of  in  bromides, 
and  when  in  the  presence  of 
chlorine,  experiment,  145. 

Bromine,  iodine,  fluorine,  problems  in 
connection  with,  149. 

Bromine,  occurrence  and  preparation, 
133. 

Bromine,  physical  properties  of,  135. 

Bromine,  preparation  of  from  potas- 
sium bromide,  manganese  dioxide, 
and  sulphuric  acid,  experiment, 
143. 

Bromine,  solidification  of,  experi- 
ment, 143. 

Bromine  water,  preparation  of,  experi- 
ment, 143. 

Bronzes,  297,  310. 

Bunsen  burner,  201. 

Bunsen  flame,  effect  of,  cooling  the, 
experiment,  217.  ) 

Burning  in  oxygen,  8. 

Burning  in  oxygen,  explanation  of,  8. 

Butter  of  tin,  311. 

Cadmium,  260. 

Cadmium  sulphide,  precipitation  of, 

experiments,  264. 
Caesium,  232. 
Calamine,  258. 
Calcite,  243. 
Calcium  carbide,  224. 
Calcium  carbide,  experiment,  253. 
Calcium  carbonate,  243. 
Calcium  carbonate  and  acid  calcium 

carbonate,  experiment,  253. 
Calcium  carbonate,  primary  or  acid, 

243. 

Calcium  chloride,  241. 
Calcium    chloride,    preparation    and 

properties  of,  experiment,  253. 
Calcium  hydride,  240. 
Calcium  hypochlorite,  68,  241. 
Calcium,  occurrence,  preparation,  and 

properties  of,  239. 
Calcium  oxide,  240. 
Calcium,  phosphates  of,  244. 
Calcium  silicate,  246. 
Calcium,  strontium,  barium,  relations 

between,  250. 
Calcium  sulphate,  242. 
Calomel,  262. 


330 


INDEX 


Calorie,  12. 

Calorimeter,  11. 

Candle  and  oil-lamp,  199. 

Carat,  304. 

Carbide  of  aluminium,  269. 

Carbide  of  calcium,  242. 

Carbon,  189. 

Carbon,  allotropic  forms  of,  diamond, 
graphite,  189. 

Carbon,  amorphous  forms  of,  190. 

Carbonates,  322. 

Carbonates  formed  from  silicates,  206. 

Carbonates,  hydrolysis  of,  225. 

Carbonates,  hydrolyzed  by  water,  196. 

Carbonation,  206. 

Carbon  bisulphide,  198. 

Carbon  bisulphide,  properties  of,  ex- 
periment, 215. 

Carbon,  comparison  of  the  different 
forms  of,  experiment,  209. 

Carbon,  compounds  of  with  hydrogen, 
192. 

Carbon,  compounds  of  with  oxygen, 
193. 

Carbon,  different  forms  contain  differ- 
ent amounts  of  energy,  191. 

Carbon  dioxide,  194. 

Carbon  dioxide,  chemical  properties 
of,  195. 

Carbon  dioxide  formed  by  burning 
carbon  in  oxygen,  experiment,  212. 

Carbon  dioxide  formed  by  the  action 
of  acids  on  carbonates,  experiment, 
213. 

Carbon  dioxide  formed  when  magne- 
site  is  heated,  experiment,  212. 

Carbon  dioxide  heavier  than  air,  ex- 
periment, 214. 

Carbon  dioxide  in  the  air,  in  large 
quantities  in  air  exhaled  from  the 
lungs,  experiment,  212. 

Carbon  dioxide  in  the  atmosphere,  108. 

Carbon  dioxide  in  the  atmosphere,  de- 
tection of,  experiment,  113. 

Carbon  dioxide,  physical  properties 
of,  197. 

Carbon  dioxide,  preparation  of,  195. 

Carbon  dioxide,  preparation  on  the 
large  scale,  experiment,  213. 

Carbon  dioxide,  reduction  by  plants, 
196. 


Carbon  dioxide,  solid,  experiment,  214. 

Carbon  dioxide  will  not  burn  and  will 
not  support  combustion,  experi- 
ment, 214. 

Carbon  monoxide,  193. 

Carbon  monoxide,  preparation  of,  ex- 
periment, 211. 

Carbon  monoxide,  reducing  power  of, 
experiment,  212. 

Carbon  monoxide,  thermochemistry 
of,  194. 

Carbon,  physical  properties  of,  191. 

Carbon,  reducing    power   of,   experi- 
ment, 210. 

Carbon,  role  of  in  producing  light, 
199. 

Carborundum,  207. 

Carnallite,  228,  256,  257. 

Carre  ice  machines,  79. 

Cassiterite,  310. 

Cast-iron,  275. 

Cast-iron,  chilled,  275. 

Cast-iron,  gray,  275. 

Cast-iron,  white,  275. 

Catalytic  reaction,  138,  156. 

Cathode,  266. 

Cations,  64. 

Celestite,  248. 

Cements,  268. 

Cements,  hydraulic,  268. 

Cerium,  207. 

Cerussite,  309. 

Chalcopyrite,  296. 

Chalk,  239,  243. 

Chamber  acid,  158. 

Charcoal,  190. 

Charcoal  absorbs  gases  and  coloring 
matter,  experiment,  210. 

Chemical  action,  laws  of,  50. 

Chemical  combination,  6. 

Chemical  elements,  4. 

Chemical  or  intrinsic  energy,  53. 

Chemical  properties  and  atomic 
weights,  combining  power,  129. 

Chemical  reaction  changes  the  compo- 
sition of  substances,  2. 

Chemistry  a  branch  of  natural 
science,  1. 

Chemistry,  kind  of  phenomena  with 
which  it  is  concerned,  1. 

Chemistry,  science  of,  1. 


INDEX 


331 


Chili  saltpetre,  220,  223. 

Chlorates,  70. 

Chlorates,  ion  of,  and  the  chlorine 
ion,  70. 

Chloric  acid,  69. 

Chloric  acid,  oxidizing  power  of, 
experiment,  75. 

Chloric  acid,  properties  of,  70. 

Chlorine,  action  on  hydrogen,  44. 

Chlorine,  action  on  organic  compounds, 
45. 

Chlorine,  action  on  organic  compounds, 
experiment,  49. 

Chlorine,  action  on  water,  45. 

Chlorine,  bleaching  action  of,  experi- 
ment, 49. 

Chlorine,  burning  hydrogen  in,  experi- 
ment, 49. 

Chlorine,  certain  physical  properties 
of,  46. 

Chlorine,  chemical  properties  of,  44. 

Chlorine,  combination  with  the  metals, 
experiment,  49. 

Chlorine,  combustion  in,  44. 

Chlorine,  comparative  inactivity  of 
dry,  47. 

Chlorine,  compounds  with  oxygen,  67. 

Chlorine,  compounds  with  hydrogen 
and  oxygen,  62,  68. 

Chlorine,  Deacon  process  of  making, 
43. 

Chlorine,  experiments  with,  47. 

Chlorine,  experiments  with  the  com- 
pounds of,  with  hydrogen  and  oxy- 
gen, 73. 

Chlorine  hydrate,  46. 

Chlorine  ion,  and  the  ion  of  chlorates, 
70. 

Chlorine  made  from  bleaching-powder 
and  electrolytically,  44. 

Chlorine,  occurrence  and  preparation 
of,  43. 

Chlorine,  preparation  from  hydro- 
chloric acid  and  manganese  di- 
oxide, experiment,  47. 

Chlorine,  preparation  from  hydro- 
chloric acid  and  potassium  di- 
chromate,  experiment,  48. 

Chlorine,  preparation  from  sodium 
vchloride,  manganese  dioxide,  and 
sulphuric  acid,  experiment,  48. 


Chlorine,  problems  in  connection  with, 
124. 

Chlorine,  volume  relations  in  which  it 
combines  with  hydrogen,  62. 

Chlorine  water,  45. 

Chlorine  water,  action  of  light  on,  ex-x 
periment,  49. 

Chromates,  290. 

Chromates  and  dichromates,  experi- 
ment, 293. 

Chromates  and  dichromates,  oxidizing 
power  of,  experiment,  294. 

Chromic  acid,  289. 

Chromic  acid,  per-,  289. 

Chromic  chromium,  action  of  reagents 
on,  experiments,  294. 

Chromic  oxide,  288. 

Chromite  of  iron,  288. 

Chromium,  288. 

Chromium  ions,  valence  and  properties 
of,  289. 

Chromium,  oxides  of,  288. 

Chromium  trioxide,  288. 

Chromium  trioxide,  formation  of,  ex- 
periment, 293. 

Chromous  oxide,  288. 

Cinnabar,  261. 

Clay,  266,  270. 

Cobalt,  280. 

Cobalt  chloride,  287. 

Cobaltic  compounds,  281. 

Cobaltic  oxide,  281. 

Cobaltite,  280. 

Cobaltous  compounds,  281. 

Cobaltous  oxide,  281. 

Cobalt  salts,  action  of  hydrogen  sul- 
phide on,  experiment,  284. 

Coins,  297. 

Color,  change  in  with  change  in  elec- 
trical charge,  280. 

Combination,  chemical,  6. 

Combining  numbers  and  atomic 
weights,  115. 

Combining  numbers,  chemical  meth- 
ods of  determining,  115. 

Combining  power  of  the  elements, 
129. 

Combustion,  9. 

Combustion  consists  in  union  with 
oxygen, 10. 

Combustion,  increase  in  weight  in,  9. 


332 


INDEX 


Combustion,  measurement  of  the  heat 

of,  11. 

Combustion,  oxygen  used  up  in,  9. 
Composition  of  substances  changed  in 

chemical  reactions,  2. 
Compounds  and  elements,  3. 
Compounds  of  the   metals,  relations 

between  the,  322. 
Conservation  of  energy,  52. 
Conservation  of    energy,  importance 

of  for  the  science  of   chemistry, 

53. 

Conservation  of  mass,  law  of,  51. 
Constant  proportion,  law  of,  51. 
Coal,  or  stone  coal,  190." 
Copper,  action  of  reagents  on  salts  of, 

experiment,  304. 
Copper,  alloys  of,  297. 
Copper,  occurrence    and  preparation 

of,  296. 

Copper,  oxides  of,  297. 
Copper,  precipitated  by  zinc  and  iron, 

298. 
Copper,  precipitated  from  its  salts  by 

certain  metals,  experiment,  305. 
Copper    precipitates    certain    metals 

from  their  salts,  experiment,  305. 
Copper,  properties  of,  296. 
Copper,  purified  by  electrolysis,  296. 
Copper  pyrites,  150. 
Copper  sulphate,  298. 
Correlation  of  energy,  52. 
Corrosive  sublimate,  262. 
Corundum,  266,  267. 
Cry  of  tin,  360. 
Cryolite,  139,  221,  266,  268. 
Crystallization,  fractional,  136. 
Cupellation,  299. 
Cupric  compounds,  297. 
Cupric  oxide,  297. 
Cuprite,  296. 
Cuprous  compounds,  297. 
Cuprous  oxide,  297. 
Cyanide  process,  303. 
Cyanogen,  198. 

Cyanogen,    preparation    of,     experi- 
ment, 216. 

Daguerre,  302. 

Davy  lamp,  203,  316. 

Deacon  process  for  making  chlorine,  43. 


Deliquescence,  38. 

Deliquescence  of  calcium  chloride,  ex- 
periment, 42. 

Deliquescent  salt,  223. 

Detonating  gas,  21. 

Developer,  301. 

Diacid  base,  99. 

Diamond,  189. 

Dibasic  acid,  99. 

Dichromates,  290. 

Dichromates  and  chromates,  experi- 
ment, 293. 

Dichromates  and  chromates,  oxidizing 
power  of,  experiment,  294. 

Diffusion  of  gases,  law  of,  25. 

Dimorphous  substances,  151. 

Disinfectant,  45. 

Disinfectant,  mercuric  chloride  a,  262. 

Disinfecting  action  of  hydrogen  di- 
oxide, 61. 

Dissociation,  hydrolytic,  176. 

Dissociation,  importance  of  the  theory 
of  electrolytic,  64. 

Dissociation,  theory  of  electrolytic,  63. 

Distillation,  34. 

Distillation  of  water,  experiment,  41. 

Dobereiner  lamp,  317. 

Dolomite,  239,  256. 

Drummond  light,  23,  240. 

Dumas'  method  for  determining  vapor 
density,  117. 

Earth,  composition  of,  6. 

Earthenware,  271. 

Efflorescence,  38. 

Efflorescence  of  sodium  sulphate,  ex- 
periment, 42. 

Efflorescent  salt,  224. 

Electric  furnace,  189. 

Electric  light,  204. 

Electrolysis  made  use  of  in  purifying 
copper,  296. 

Electrolysis  of  water,  35. 

Electrolytic  dissociation,  importance 
of  the  theory  of,  64. 

Electrolytic  dissociation,  theory  of,  63. 

Electrolytic  gas,  22. 

Electroplating,  282. 

Elements  and  compounds,  3. 

Elements  and  compounds,  number 
of,  3. 


INDEX 


333 


Elements,  chemical,  4. 
Endothermic  reactions,  81.    ' 
Energy,  chemical  or  intrinsic,  53. 
Energy,  correlation  and  conservation 

of,  52. 
Energy,  importance  of  the  conservation 

of,  for  the  science  of  chemistry,  53. 
Epsom  salt,  257. 

Equations  and  formulas,  chemical,  54. 
Etching  glass,  experiment,  148. 
Ether,  198. 
Ethylene,  192. 
Exothermic  reactions,  81. 

Faraday's  law  the  basis  of  chemical 
valence,  72,  277. 

Feldspar,  66,  227,  270. 

Ferric  and  ferrous  compounds,  277. 

Ferric  chloride,  278. 

Ferric  chlorine,  preparation  of,  exper- 
iment, 283. 

Ferric  hydroxide,  278. 

Ferric  sulphate,  279. 

Ferricyanide  of  potassium,  280. 

Ferrocyanide  of  potassium,  279. 

Ferrous  and  ferric  compounds,  277. 

Ferrous  and  ferric  salts,  detection  of, 
experiment,  283. 

Ferrous  chloride,  278. 

Ferrous  chloride,  preparation  of,  ex- 
periment, 282. 

Ferrous  hydroxide,  278. 

Ferrous  salts,  properties  of,  experi- 
ment, 283. 

Ferrous  sulphate,  279. 

Fertilizer,  commercial,  245. 

Filter-paper,  33. 

Filtration,  33. 

Fire-brick,  270. 

"Fixing"  photographs,  161,  301. 

Flames  and  their  luminosity,  200. 

Flames  differ  greatly  in  their  luminos- 
ity, experiment,  216. 

Flames,  effect  of  cooling  on  luminosity, 
203. 

Flint,  205. 

Fluorine,  139. 

Fluorine,  chemical  properties  of,  141. 

Fluorine,  liquefaction  of,  141. 

Fluorine,  occurrence  and  preparation, 


Fluorine,  physical  properties  of,  141. 

Fluor-spar,  139,  239. 

Flux,  226. 

Formulas  and  equations,  chemical,  54. 

Fractional  crystallization,  136. 

Fuming  nitric  acid,  85. 

Fuming  sulphuric  acid,  156. 

Galena,  150,  307,  308. 

Gallium,  271. 

Galvanized  objects,  258. 

Garnet,  246. 

Garnets,  270. 

Garnierite,  282. 

Gases,  problems  in  connection  with 
the  volumes  of,  125. 

Gasometer,  166. 

Gay-Lussac's  law  for  gases,  112. 

Gay-Lussac  tower,  158. 

Generalization,  50. 

Germanium,  207. 

German  silver,  258,  282,  297. 

Gilchrist-Thomas  converter,  276. 

Glass,  246. 

Glass,  Bohemian  or  potassium,  246. 

Glasses,  colored,  246. 

Glass,  flint,  246. 

Glass,  soda,  246. 

Glauber's  salt,  223. 

Glazing,  271. 

Glover  tower,  157. 

Glucinum,  256. 

Gneisses,  205. 

Gold,  302. 

Gold,  metallurgy  of,  303. 

Gold,  oxides  and  hydroxides  of,  304. 

Gold-plating,  304. 

Gold,  properties  of,  303. 

Gold,  salts  of,  304. 

Goldschmidt's  mixtures,  285,  288. 

Granites,  205. 

Graphite,  or  plumbago,  189. 

Green  vitriol,  279. 

Group  I  in  qualitative  analysis,  324. 

Group  II  in  qualitative  analysis,  324. 

Group  III  in  qualitative  analysis,  324. 

Group  IV  in  qualitative  analysis,  324. 

Group  V  in  qualitative  analysis,  325. 

Group  VI  in  qualitative  analysis,  325. 

Group  I  of  the  Periodic  System,  rela- 
tions within,  318. 


334 


INDEX 


Group  II  of  the  Periodic  System,  rela- 
tions within,  319. 

Group  III  of  the  Periodic  System,  re- 
lations within,  320. 

Group  IV  of  the  Periodic  System,  rela- 
tions within,  320. 

Group  V  of  the  Periodic  System,  rela- 
tions within,  320. 

Group  VI  of  the  Periodic  System,  re- 
lations within,  321. 

Group  VII  of  the  Periodic  System,  re- 
lations within,  321. 

Group  VIII  of  the  Periodic  System, 
relations  within,  321. 

Gun  metal,  297. 

Gunpowder,  230. 

Gypsum,  239,  242. 

Gypsum,  plaster  of  paris,  experiment, 
253. 

Halides,  322, 

Halogen  acids,  comparisons  of  the 
several,  142. 

Hard  water,  32,  245. 

Heat  evolved  in  the  reaction  and 
stability  of  the  product,  38. 

Heat  of  combustion,  measurement  of, 
11. 

Heat  of  formation  and  decomposition, 
12. 

Heat  of  fusion  of  ice,  39. 

Heat  of  neutralization,  101. 

Heat  of  neutralization  constant,  ex- 
planation of,  102. 

Heat  of  neutralization, experiment,  104. 

Heat  of  vaporization  of  water,  39. 

Heavy  spar,  248,  249. 

Helium,  111. 

Hematite,  273,  278. 

Homologous  series  of  compounds,  193. 

Hornblende,  256. 

Hydraulic  cements,  268. 

Hydraulic  mining,  303. 

Hydride  of  silicon,  205. 

Hyldride  of  sodium,  221. 

Hydriodic  acid,  138. 

Hydrobromic  acid,  135. 

Hydrobromic  acid,  preparation  of,  by 
the  action  of  sulphuric  acid  on 
potassium  bromide,  experiment, 
144. 


Hydrobromic  acid,  properties  of,  135, 

145. 

Hydrochloric  acid,  62. 
Hydrochloric  acid,   action  on  metals 

like  zinc,  65. 
Hydrochloric  acid,   aqueous  solution 

of,  67. 

Hydrochloric  acid,  chemical  proper- 
ties of,  63. 

Hydrochloric  acid,  detection  of,  66. 
Hydrochloric  acid  from  sulphuric  acid 

on  concentrated  hydrochloric  acid, 

experiment,  73. 
Hydrochloric  acid,  physical  properties 

of,  67. 

Hydrochloric  acid,  preparation  of,  62. 
Hydrochloric    acid,  prepared  by  the 

action  of  sulphuric  acid  on  sodium 

chloride,  experiment,  73. 
Hydrochloric  acid,   properties  of  an 

aqueous  solution  of,  experiment,  74. 
Hydrochlorplatinic  acid,  318. 
Hydrocyanic  acid,  199. 
Hydrofluoric  acid,  142. 
Hydrofluoric    acid,    preparation    of, 

etching  glass,  experiment,  148. 
Hydrogen,  19. 
Hydrogen  and  oxygen,  relations  by 

volume  in  which  they  combine,  22. 
Hydrogen,  combination  with  oxygen, 

20. 

Hydrogen,  combustion  of,  27. 
Hydrogen,  diffusion  of,  experiment,  31. 
Hydrogen  dioxide,  60. 
Hydrogen  dioxide  a  disinfectant,  61. 
Hydrogen  dioxide,  properties  of,  60. 
Hydrogen,  equations  involved  in  the 

study  of,  56. 

Hydrogen,  experiments  with,  26. 
Hydrogen  lighter  than  air,  29. 
Hydrogen,  liquefaction  of,  25. 
Hydrogen,  nascent,  24. 
Hydrogen,  physical  properties  of,  24. 
Hydrogen,  preparation  of,  19. 
Hydrogen,    preparation     of,    experi- 
ment, 26. 
Hydrogen,    problems    in    connection 

with,  123. 

Hydrogen,  properties  of  liquid,  26. 
Hydrogen,  rate  of  diffusion,  25. 
Hydrogen,  reducing  power  of,  23. 


INDEX 


335 


Hydrogen,  reducing  power  of,  experi- 
ment, 28. 

Hydrogen  soap  bubbles  rise  in  the  air, 
30. 

Hydrogen  sulphide,  152. 

Hydrogen  sulphide,  action  of  on  salts 
of  metals,  experiment,  165. 

Hydrogen  sulphide  burns  in  contact 
with  air,  experiment,  165. 

Hydrogen  sulphide,  chemical  proper- 
ties of,  153. 

Hydrogen  sulphide,  dissociation  of, 
154. 

Hydrogen  sulphide  group,  153. 

Hydrogen  sulphide,  physical  proper- 
ties of,  154. 

Hydrogen  sulphide,  preparation  of, 
experiment,  164. 

Hydrogen  when  burned  produces 
water,  experiment,  28. 

Hydrolysis,  225. 

Hydrolytic  dissociation,  176. 

Hydrosulphides  or  acid  sulphides,  154. 

Hydroxides,  322. 

"  Hypo,"  161,  224. 

Hypobromous  acid,  136. 

Hypochlorite,  calcium,  68. 

Hypochlorous  acid,  68. 

"  Hyposulphite  of  sodium,"  161. 

Hypothesis  of  Avogadro,  116. 

Ice,  artificial  preparation  of,  79. 

Ice,  heat  of  fusion  of,  39. 

Iceland  spar,  242. 

Illumination,  199. 

Impure  or  commercial  iron,  275. 

Indicators,  103. 

Indicators,  experiments,  104. 

Insoluble  compound  formed  whenever 
possible,  71. 

Intrinsic  energy,  more  in  ozone  than 
in  oxygen,  61. 

Intrinsic  or  chemical  energy,  53. 

lodic  acid,  139. 

Iodide  of  potassium,  action  of  sul- 
phuric acid  on,  147. 

Iodine,  137. 

Iodine,  chemical  properties  of,  138. 

Iodine,  compounds  with  oxygen  and 
hydrogen,  139. 

Iodine,  detection  of,  147. 


Iodine,  detection  of  in  the  presence  of 
chlorine  and  bromine,  experiment, 
147. 

Iodine,  occurrence  and  preparation, 
137. 

Iodine,  physical  properties  of,  138. 

Iodine,  preparation  of  from  potassium 
iodide,  manganese  dioxide,  and  sul- 
phuric acid,  experiment,  146. 

Iodine,  properties  of,  action  on  phos- 
phorus and  zinc,  experiment,  146. 

Iodine,  properties  of,  sublimation,  solu- 
bility reaction  with  starch,  experi- 
ment, 146. 

Iodine,  tincture  of,  138. 

Ions,  method  of  indicating  the  exist- 
ence of ,  65. 

Ions,  nature  and  role  of  in  chemistry, 
325. 

Iridium,  315. 

Iron  alums,  279. 

Iron,  occurrence  and  preparation,  273. 

Iron,  precipitation  of  as  the  sulphide, 
experiment,  284. 

Iron,  properties  of,  274. 

Iron,  sulphides  of,  278. 

Iron  vitriol,  279. 

Isomorphous  substances,  269. 

Jasper,  205. 

Kainite,  231. 
Kaolin,  266, 270. 
Kieserite,  256,  257. 
Kindling  temperature,  203. 
Krypton,  111. 

Lamp-black,  or  soot,  190. 

Lanthanum,  271. 

Laughing  gas,  81. 

Laws  of  chemical  action,  50. 

Lead,  307. 

Lead  acetate,  309. 

Lead  carbonate,  308. 

Lead  chloride,  experiment,  313. 

Lead,  chloride  of,  308. 

Lead  dioxide,  308. 

Leaden  chamber,  157. 

Lead,    occurrence,    preparation,    and 

properties,  307. 
Lead,  oxides  of,  308. 


836 


INDEX 


Lead,  oxides  of,  experiment,  312. 
Lead,  precipitated  from  its  salts  by 

metals,  308. 
Lead  precipitated  from  its  salts  by 

other  metals,  experiment,  312. 
Lead  sesquioxide,  308. 
Lead  suboxide,  308. 
Lead,  sugar  of,  309. 
Lead  sulphide,  308. 
Lead  sulphide,  experiment,  313. 
Lead  tree,  312. 
Lead,  white,  309. 
Le  Blanc  method  for  preparing  sodium 

•  carbonate,  224. 
Leclanche  cell,  286. 
Light,  action  on  silver  chloride  and 

bromide,  experiment,  306. 
Lignite,  191. 
Lime,  240. 
Lime    formed    by    heating    calcium 

carbonate,  experiment,  252. 
Lime,  slaking  of,  preparation  of  cal- 
cium hydroxide,  experiment,  252. 
Limestone,  239,  243. 
Lime-water,  240. 
Limonite,  273. 
Liquid  air,  109. 
Litharge,  308. 
Lithium,  232. 
Litmus,  103. 
Luminosity,  effect  of  an  indifferent 

gas  on,  201. 

Luminosity,  effect  of  pressure  on,  201. 
Luminous  flame  from  a  non-luminous 

flame,  experiment,  216. 

Magnesia,  257. 

Magnesite,  256. 

Magnesium,  256. 

Magnesium  burns  in  the  air,  experi- 
ment, 263. 

Magnesium  chloride,  257. 

Magnesium  nitride,  257. 

Magnesium  oxide  and  hydroxide,  257. 

Magnesium  phosphate,  formation  of, 
experiment,  264. 

Magnesium,  phosphates  of,  257. 

Magnesium  sulphate,  257. 

Magnetite,  273,  278. 

Malachite,  296. 

Manganate  of  potassium,  287. 


Manganate  of  potassium,  preparation 
of,  experiment,  291. 

Manganese,  285. 

Manganese,  occurrence,  preparation, 
properties,  285. 

Manganese,  oxides  of,  285. 

Manganese,  quadrivalent,  286. 

Manganese  salts,  reaction  of,  experi- 
ment, 291. 

Manganese,  valence  of,  286. 

Manganic  acid,  287. 

Manganic  acid,  per-,  287. 

Manganic  compounds,  286. 

Manganous  salts,  286. 

Mantle  of  Welsbach  light,  203. 

Marble,  239,  243. 

Marl,  270. 

Marsh's  method  for  detecting  arsenic, 
178. 

Mass  action,  law  of,  250. 

Mass,  effect  of  on  chemical  activity, 
207. 

Massicot,  308. 

Mass,  law  of  the  conservation  of,  51. 

Matches,  safety,  230. 

Mercuric  chloride  a  disinfectant,  262. 

Mercuric  sulphide,  263. 

Mercurous  and  mercuric  chlorides,  262. 

Mercurous  and  mercuric  oxides,  261. 

Mercury,  261. 

Mercury,  properties  of,  261. 

Mercury  sulphide,  precipitation  of, 
experiment,  264. 

Metals,  219. 

Metals,  relations  between  the  com- 
pounds of,  322. 

Metals,  the  chemistry  of  the  chem- 
istry of  the  metal  ions,  219. 

Metaphosphate  formed  by  heating  mi- 
crocosmic  salt,  experiment,  185. 

Metaphosphoric  acid,  176. 

Metastannic  acid,  311. 

Methane,  192. 

Methane,  preparation  of,  experiment, 
211. 

Methyl  orange,  103. 

Mica,  246,  266,  270. 

Microcosmic  salt,  234. 

Milk  of  lime,  240. 

Mineral  waters,  33. 

Mining  hydraulic,  303. 


INDEX 


337 


Mining,  placer,  303. 

Minium,  or  red  lead,  308. 

Mirrors,  300. 

Moissan's  method  of  preparing  fluo- 
rine, 140. 

Molecular  weights,  atomic  weights 
from,  119. 

Molecule,  52. 

Molybdenum,  290. 

Monacid  base,  99. 

Monobasic  acid,  99. 

Multiple  proportions,  law  of,  51. 

Nascent  hydrogen,  24. 

Natural  science,  chemistry  a  branch 
of,  1. 

Needle  valve,  110. 

Neodymium,  172. 

Neon,  111. 

Neutralization,  explanation  of  the 
constant  heat  of,  102. 

Neutralization,  heat  of,  101. 

Neutralization  of  acids  and  bases,  98. 

Neutralization  of  acids  by  bases,  ex- 
periment, 104. 

Neutralization  of  acids  by  bases,  quan- 
titative work,  experiment,  105. 

Niccolite,  282. 

Nickel,  282. 

"Nickel  coin,"  282. 

Nickel,  compounds  of,  282. 

Nickel  salts,  action  of  hydrogen  sul- 
phide on,  experiment,  284. 

Nitrates,  322. 

Nitrates,  dissociation  of,  85. 

Nitric  acid,  83. 

Nitric  acid,  acid  properties  of,  experi- 
ment, 95. 

Nitric  acid,  action  of  on  hydrochloric 
acid,  experiment,  97. 

Nitric  acid,  chemical  properties  of,  84. 

Nitric  acid,  detection  of,  84. 

Nitric  acid,  detection  of,  experiment, 
97. 

Nitric  acid,  dissociation  of,  85. 

Nitric  acid  dissolves  certain  metals, 
experiment,  97. 

Nitric  acid,  fuming,  85. 

Nitric  acid,  fuming,  experiment,  96. 

Nitric  acid,  oxidizing  power  of,  ex- 
periments, 96. 


Nitric  acid,  preparation  of,  experi- 
ment, 94. 

Nitric  oxide,  81. 

Nitric  oxide,  combination  of  with  oxy- 
gen, experiment,  93. 

Nitric  oxide,  does  not  support  com- 
bustion, experiment,  93. 

Nitric  oxide,  preparation  of,  experi- 
ment, 92. 

Nitride  of  magnesium,  257. 

Nitrifying  ferment,  230. 

Nitrogen,  76. 

Nitrogen,  compound  with  hydrogen, 
ammonia,  77. 

Nitrogen,  compounds  with  hydrogen 
and  oxygen,  experiment,  86. 

Nitrogen,  compounds  with  oxygen,  80. 

Nitrogen ,  compounds  with  oxygen  and 
hydrogen,  80. 

Nitrogen,  dioxide  or  peroxide,  82. 

Nitrogen,  in  the  atmosphere,  108. 

Nitrogen  obtained  from  the  air,  ex- 
periment, 86. 

Nitrogen  pentoxide,  82. 

Nitrogen  peroxide,  preparation  of, 
experiment,  94. 

Nitrogen,  preparation  from  ammo- 
nium nitrite,  experiment,  86. 

Nitrogen,  problems  in  connection  with, 
124. 

Nitrogen,  properties  of,  76. 

Nitrogen,  sesquioxide  or  trioxide,  82. 

Nitrogen,  sesquioxide  or  trioxide, 
preparation  of,  experiment,  93. 

Nitrosulphonic  acid,  158. 

Nitrosyl-sulphuric  acid,  158. 

Nitrous  acid,  82. 

Nitrous  oxide,  80. 

Nitrous  oxide,  preparation  of,  experi- 
ment, 92. 

Nitrous  oxide  supports  combustion, 
experiment,  92. 

Noble  metals,  298. 

Number  of  elements  and  compounds,  3. 

Octivalent  ion,  72. 
Onyx,  205. 
Opal,  205. 
Organic  acids,  198. 
Orthoclase,  270. 
Orthophosphoric  acid,  175. 


338 


INDEX 


Osmium,  315. 

Oxidation  in  terms  of  Faraday's  law, 
278. 

Oxidation,  rapid  and  slow,  11. 

Oxides,  322. 

Oxidizing  flame,  202. 

Oxygen,  7. 

Oxygen  and  hydrogen  combine  with 
li Deration  of  heat,  22. 

Oxygen  and  hydrogen,  relations  by 
volume  in  which  they  combine, 
22. 

Oxygen,  certain  physical  properties 
of,  12. 

Oxygen,  combination  with  hydrogen, 
20. 

Oxygen,  combustion  of  carbon  in,  ex- 
periment, 16. 

Oxygen,  combustion  of  iron  in,  ex- 
periment, 17. 

Oxygen,  combustion  of  phosphorus  in, 
experiment,  17. 

Oxygen,  combustion  of  sulphur  in,  ex- 
periment, 16. 

Oxygen,  equations  involved  in  the 
study  of,  55. 

Oxygen,  experiments  with,  13. 

Oxygen,  explanation  of  burning  in,  8. 

Oxygen  in  the  atmosphere,  determina- 
tion of  the  amount  of,  experiment, 
111. 

Oxygen,  liquefaction  of,  13. 

Oxygen,  occurrence  of,  7. 

Oxygen,  preparation  from  a  mixture 
of  manganese  or  lead  dioxide  and 
hydrogen  dioxide,  experiment,  16. 

Oxygen,  preparation  from  a  mixture 
of  potassium  chlorate  and  manga- 
nese dioxide,  experiment,  15. 

Oxygen,  preparation  from  manganese 
dioxide,  experiment,  14. 

Oxygen,  preparation  from  mercuric 
oxide,  experiment,  15. 

Oxygen,  preparation  from  potassium 
chlorate,  experiment,  13. 

Oxygen,  preparation  of,  7. 

Oxygen,  pressure  of,  varies  with  the 
conditions,  12. 

Oxygen,  problems  in  connection  with, 
123. 

Oxygen,  substances  burn  readily  in,  8. 


Oxygen  used  up  in  combustion,  expei 

iment,  18. 

Oxy hydrogen  blowpipe,  22. 
Ozone,  allotropic  modification  of  oxj 

gen,  58. 

Ozone  and  hydrogen  dioxide,  58. 
Ozone  and  oxygen,  same  kinds  of  ma 

ter,  but  different  amounts  of  er 

ergy,  59. 
Ozone  contains  more  intrinsic  energ 

than  oxygen,  60. 
Ozone,  preparation  of,  58. 
Ozone,  transformation  into  oxygen,  5! 

Palladium,  315. 

Palladium  hydride,  315. 

Paris,  plaster  of,  242. 

Passive  condition,  278. 

Peat,  191. 

Perchloric  acid,  70. 

Perchloric  acid,  properties  of,  71. 

Perchromic  acid,  289. 

Periodic  acid,  139. 

Periodic  system,  126. 

Permanganate  of  potassium,  288. 

Permanganate  of  potassium,  prepan 

tion  of,  experiment,  292. 
Permanganic  acid,  287. 
Phenolphthalein,  103.    " 
Phenomena,    the    kind    with    whic 

chemistry  is  concerned,  1. 
Philosopher's  wool,  258. 
Phosphate  beds,  172. 
Phosphates,  324. 
Phosphates,  normal,  175. 
Phosphates,  primary,  175. 
Phosphates,  secondary,  175. 
Phosphine,  174. 
Phosphorescent,  172. 
Phosphoric  acid,  available,  246. 
Phosphoric  acid,  dissociation  of,  175. 
Phosphoric  acid,  formed  when  wate 

acts     on    phosphorus    pentoxid( 

experiment,  184. 
Phosphoric  acid,  insoluble,  246. 
Phosphoric  acid,meta-,  176. 
Phosphoric  acid,   or  orthophosphor; 

acid,  175. 

Phosphoric  acid,  pyro-,  176. 
Phosphorite,  172,  245. 
Phosphorous  acid,  176. 


INDEX 


339 


Phosphorus,  172. 

Phosphorus,  acids  of,  174. 

Phosphorus,  an  active  element  chemi- 
cally, experiment,  182. 

Phosphorus,  arsenic,  antimony,  bis- 
muth, problems  in  connection 
with,  188. 

Phosphorus,  compounds  with  hydro- 
gen, experiment,  183. 

Phosphorus,  compounds  with  oxygen 
and  hydrogen,  174. 

Phosphorus,  occurrence  and  prepara- 
tion, 172. 

Phosphorus  pentachloride,  177. 

Phosphorus,  properties  of,  172. 

Phosphorus,  red,  173. 

Phosphorus,  red,  more  stable  than 
yellow,  experiment,  183. 

Phosphorus  trichloride,  177. 

Phosphorus,  yellow,  173. 

Photochemical  reactions,  45. 

Photography,  301. 

Pig-iron,  274,  275. 

Placer  mining,  303. 

Plaster  of  paris,  242. 

Platinic  acid,  hydrochlor-,  318. 

Platinum,  315. 

Platinum  black,  316. 

Platinum,  chlorides  of,  318. 

Platinum ,  oxid  es  and  hydroxides  of ,  318. 

Platinum,  properties  of,  316. 

Platinum  sponge,  316. 

Platinum,  uses  of,  317. 

Plumbago,  or  graphite,  189. 

Polymorphous  substances,  151. 

Porcelain,  271. 

Positive  photograph,  302. 

Potash,  caustic,  228. 

Potassium,  227. 

Potassium  acid  carbonate,  232. 

Potassium  acid  sulphate  reacts  acid, 
231. 

Potassium  carbonate,  232. 

Potassium  chlorate,  229. 

Potassium  chlorate,  preparation  of, 
experiment,  74. 

Potassium  chlorate,  quantitative  de- 
termination of  oxygen  in,  experi- 
ment, 236. 

Potassium  chloride,  228. 

Potassium  ferricyanide,  280. 


Potassium  ferrocyanide,  279. 

Potassium  hydride,  228. 

Potassium  hydroxide,  228. 

Potassium  hydroxide  precipitates  the 
hydroxides  of  the  heavy  metals, 
experiment,  236. 

Potassium  manganate,  287. 

Potassium  manganate,  preparation  of, 
experiment,  291. 

Potassium  nitrate,  230. 

Potassium  nitrate  mixed  with  carbon 
burns  readily,  experiment,  237. 

Potassium  permanganate,  288. 

Potassium  permanganate,  color  of, 
experiment,  292. 

Potassium  peroxide,  228. 

Potassium,  phosphates  of,  232. 

Potassium,  properties  of,  227. 

Potassium  sulphate,  231. 

Precious  metals,  298. 

Prints,  302. 

Prismatic  colors,  251. 

Problems  in  connection  with  bromine, 
iodine,  and  fluorine,  149. 

Problems  in  connection  with  calcium, 
strontium,  and  barium,  255. 

Problems  in  connection  with  carbon, 
silicon,  and  boron,  208. 

Problems  in  connection  with  iron,  co- 
balt, and  nickel,  284. 

Problems  in  connection  with  magne- 
sium, zinc,  cadmium,  and  mercury, 
265. 

Problems  in  connection  with  phos- 
phorus, arsenic,  antimony,  and  bis- 
muth, 188. 

Problems  in  connection  with  sodium 
and  potassium,  238. 

Problems  in  connection  with  sulphur, 
selenium,  and  tellurium,  170. 

Problems  with  lead  and  tin,  314. 

Properties  of  the  elements  and  atomic 
weights,  relations  between,  126. 

Proportion,  law  of  constant,  51. 

Proportions,  law  of  multiple,  51. 

Praseodymium,  172. 

Pyrites,  150,  273. 

Pyrolusite,  285. 

Pyrophosphate  formed  by  heating  a  sec- 
ondary phosphate,  experiment,  184. 

Pyrophosphoric  acid,  176. 


340 


INDEX 


Quadrivalent  ion,  72. 

Quantitative  method  of  dealing  with 

chemical  reactions,  122. 
Quartation,  303. 
Quartz,  204,  205. 
Quinquivalent  ion,  72. 

Radiation  of  uranium,  291. 

Radium,  polonium,  thorium,  291. 

Reaction  flame  of  calcium,  experi- 
ment, 253. 

Reactions  chemical,  quantitative  meth- 
od of  dealing  with,  122. 

Red  lead,  308. 

Red  prussiate  of  potash,  280. 

Reducing  agent,  24. 

Reducing  flame,  202. 

Reduction  by  hydrogen,  23. 

Reduction  in  terms  of  Faraday's  law, 
278. 

Regia,  aqua,  85. 

Relations  between  the  compounds  of 
the  metals,  322. 

Relations,  general,  319. 

Relations  within  the  groups  of  the 
Periodic  System,  319. 

Rhodium,  315. 

Rock  crystal,  205. 

Rose's  fusible  metal,  181. 

Rubidium,  232. 

Ruby,  266,  267. 

Rupert,  Prince,  drops,  247. 

Rusting  of  iron,  275. 

Ruthenium,  315. 

Safety  lamp,  203. 

Safety  matches,  230. 

Sal  ammoniac,  233. 

Salt,  definition  of,  100. 

Saltpetre  bacteria,  230. 

Saltpetre,  Chili,  220,  223. 

Saltpetre  plantations,  230. 

Salts,  100. 

Sand,  205. 

Sandstone,  205. 

Sapphire,  266,  267. 

Saturated  and  unsaturated  solutions, 

40. 

Scandium,  271. 
Schlippe's  salt,  181. 


Science,  natural,  chemistry  a  branch 
of,  1. 

Science  of  chemistry,  1. 

Selenic  acid,  162. 

Selenide,  hydrogen,  161. 

Selenious  acid,  162. 

Selenium,  161. 

Selenium  dioxide,  161. 

Sensitive  film,  301. 

Separation  of  the  members  of  any 
group  in  qualitative  analysis,  325. 

Septivalent  ion,  72. 

Serpentine,  256. 

Sexivalent  ion,  72. 

Siderite,.  273. 

Silicates,  conversion  into  carbonates, 
206. 

Silicic  acid,  formation  of  salt  of,  ex- 
periment, 217. 

Silicic  acid,  normal,  206. 

Silicic  acids,  poly-,  206. 

Silicon,  204. 

Silicon,  acids  of,  206. 

Silicon  dioxide,  205. 

Silicon  dioxide,  preparation  of,  experi- 
ment, 217. 

Silicon  hydride,  205. 

Silicon,  preparation  of,  205. 

Silicon  tetrafluoride,  207. 

Silver,  298. 

Silver  bromide,  300. 

Silver  chloride,  300. 

Silver  chloride  and  bromide,  action  of 
light  on,  experiment,  306. 

Silvering,  299. 

Silver  iodide,  302. 

Silver  ion,  300. 

Silver  nitrate,  action  of  reagents  on, 
experiment,  305. 

Silver,  oxides  and  hydroxide  of,  300. 

Silver-plated  wares,  299. 

Silver-plating,  299. 

Silver,  precipitated  by  mercury,  ex- 
periment, 306. 

Silver,  preparation  of,  298. 

Silver,  properties  of,  299. 

Slaking  of  lime,  240. 

Smaltite,  280. 

Smithsonite,  258. 

Soda-lime,  241. 

Sodium,  220. 


INDEX 


341 


Sodium  and  potassium,  flame  reaction 
of,  experiment,  237. 

Sodium  borate,  226. 

Sodium  bromide,  223. 

Sodium  carbonate,  224. 

Sodium  carbonate,  acid,  225. 

Sodium  carbonate,  preparation  of  by 
the  ammonia  or  Solvay  process, 
experiment,  235. 

Sodium,  chemistry  of,  the  chemistry  of 
the  sodium  ion,  222. 

Sodium  chloride,  222. 

Sodium  hydride,  221. 

Sodium  hydroxide,  222. 

Sodium  hydroxide,  properties  of,  ex- 
periment, 235. 

Sodium  iodide,  223. 

Sodium  nitrate,  223. 

Sodium,  occurrence  of,  220. 

Sodium  peroxide,  221. 

Sodium  phosphate,  normal,  226. 

Sodium  phosphate,  primary,  226. 

Sodium  phosphate,  secondary,  226. 

Sodium,  phosphates  of,  226. 

Sodium,  preparation  of,  221. 

Sodium,  properties  of,  experiment,  235. 

Sodium,  properties  of,  metallic,  221. 

Sodium  pyrosulphate,  224. 

Sodium  sulphate,  223. 

Sodium  sulphate,  acid,  223. 

Sodium  thiosulphate,  224. 

Soft  solder,  310. 

Solder,  227. 

Solder,  soft,  310. 

Solutions,  unsaturated  and  saturated, 
40. 

Solvay  or  ammonia  process  for  pre- 
paring sodium  carbonate,  224. 

Soot  or  lamp-black,  190. 

Specific  heat  of  water,  40. 

Spectroscope,  251. 

Spectroscopic  analysis,  252. 

Spectroscopic  examination  of  the  alka- 
lies and  alkaline  earths,  experi- 
ment, 254. 

Spectroscopic  study  of  the  alkalies  and 
alkaline  earths,  251. 

Spiegel  iron,  275. 

Stalactites,  244. 

Stalagma,  244. 

Stalagmites,  244. 


Stannate,  sulpho-,  311. 

Stannates,  sulpho-  or  thio-,  312. 

Stannic  acid,  meta-,  311. 

Stannic  chloride,  311. 

Stannic  hydroxide,  311. 

Stannic  oxide,  310. 

Stannic  sulphide,  312. 

Stannous  chloride,  311. 

Stannous  chloride,  preparation  and 
properties  of,  experiment,  313. 

Stannous  hydroxide,  310. 

Stannous  oxide,  310. 

Stannous  sulphide,  311. 

Steel,  276. 

Stibine,  180. 

Stibine,  preparation  and  properties  of, 
experiment,  187. 

Stibnite,  150,  179. 

Still,  34. 

Stoneware,  271. 

Storage  cells,  101. 

Strontianite,  248. 

Strontium,  247. 

Strontium  flame,  reaction  of,  experi- 
ment, 254. 

Strontium  hydroxide,  248. 

Strontium  nitrate,  248. 

Strontium,  occurrence,  preparation, 
properties,  248. 

Strontium  sulphate,  248. 

Sublimation,  138 

Sugar  of  lead,  309. 

Sulphantimonious  and  sulphantimonic 
acids,  181. 

Sulphates,  323. 

Sulphates,  acid,  160. 

Sulphates,  normal,  160. 

Sulphides,  151,  324. 

Sulphides,  acid,  153. 

Sulphostannates,  312. 

Sulphur,  150. 

Sulphur,  chemical  properties  of,  151. 

Sulphur,  compounds  with  oxygen  and 
hydrogen,  155. 

Sulphur,  different  crystal  forms  of, 
experiment,  163. 

Sulphur  dioxide,  155. 

Sulphur  dioxide,  bleaching  action  of 
moist,  experiment,  168. 

Sulphur  dioxide,  liquefaction  of,  ex- 
periment, 168. 


342 


INDEX 


Sulphur  dioxide,  preparation  of  by 
burning  sulphur  in  oxygen,  ex- 
periment, 166. 

Sulphur  dioxide,  preparation  of  by 
the  action  of  sulphuric  acid  on  acid 
sodium  sulphite,  experiment,  167. 

Sulphur  dioxide,  preparation  of  by 
the  action  of  sulphuric  acid  on  cop- 
per, experiment,  167. 

Sulphur,  flowers  of,  150. 

Sulphur,  fusion  and  distillation  of, 
experiment,  162. 

Sulphuric  acid,  chemical  properties  of, 
159. 

Sulphuric  acid,  dissociation  of,  160. 

Sulphuric  acid,  fuming,  156. 

Sulphuric  acid,  physical  properties, 
160. 

Sulphuric  acid,  preparation  of,  157. 

Sulphuric  acid,  properties  of,  experi- 
ment, 170. 

Sulphuric  acid,  scientific  and  technical 
uses  of,  160. 

Sulphur,  occurrence  and  purification, 
151. 

Sulphurous  acid,  155. 

Sulphurous  acid,  dissociation  of,  156. 

Sulphurous  acid,  reducing  action  of, 
experiment,  169. 

Sulphur,  physical  properties  of,  151. 

Sulphur,  reaction  of  with  metals,  ex- 
periment, 164. 

Sulphur,  roll  or  stick,  150. 

Sulphur,  selenium,  tellurium,  prob- 
lems in  connection  with,  170. 

Sulphur  trioxide,  156. 

Sulphur  water,  152. 

Superphosphate,  245. 

Sympathetic  ink,  281. 

Sympathetic  ink,  experiment,  284. 

Synthesis  of  water,  37. 

Talc,  256. 
Tantalum,  182. 
Telluric  acid,  162. 
Telluride,  hydrogen,  162. 
Tellurious  acid,  162. 
Tellurium,  162. 
Tempering  iron,  277. 
Tetraboric  acid,  208. 
Thallium,  271. 


Theory  of  electrolytic  dissociation,  63. 
Thiosulphate  of  sodium,  161. 
Thomas-Gilchrist  converter,  276. 
Thomas  slag,  277. 
Thorium,  207. 
Tin,  309. 

Tin,  alloys  of,  310. 
Tin-butter,  311. 
Tin,  cry  of,  310. 
Tincture  of  iodine,  138. 
Tin,  hydroxides  of,  experiment,  313. 
Tin  ions,  310. 

Tin,  preparation  and  properties  of,  310. 
Tin  salt,  311. 
Tinstone,  310. 
Tin,  sulphides  of,  311. 
Tin,  sulphides  of,  experiment,  314. 
Tin  thrown  down  by  other  metals,  ex- 
periment, 313. 
Titanium,  207. 
Topaz,  oriental,  267. 
Triacid  base,  99. 
Tribasic  acid,  100. 
Trivalent  ion,  72. 
Tungsten,  290. 

Univalent  ion,  71 
Uranium,  290. 
Uranium  radiation,  291. 

Valence,  71. 

Valence,  Faraday's  law  the  basis  of, 

72. 

Valence,  variable,  262. 
Vanadium,  182. 
Vapor-density  measurements,  results 

of,  118. 

Velocity  of  reaction,  138. 
Volatile  compound  formed  whenever 

it  can  be  formed,  160. 
Volume  relations  in  which  oxygen  and 

hydrogen  combine,  22. 

Water,  boiling-point  of,  38. 

Water,  chemical  behavior  of,  37. 

Water,  composition  of,  35. 

Water,  electrolysis  of,  35. 

Water,  electrolysis  of,  experiment,  41. 

Water-gas,  194. 

Water,  hard,  32. 

Water,  heat  of  vaporization  of,  39. 


INDEX 


343 


Water  in  nature  is  impure,  32. 

Water  in  the  air,  proof  of  the  pres- 
ence of,  experiment,  114. 

Water  not  an  element  but  a  com- 
pound, 35. 

Water,  occurrence  of,  32. 

Water  of  crystallization,  or  hydra- 
tion,  32. 

Water,  physical  properties  of,  38. 

Water,  purification  of,  33. 

Waters,  mineral,  33. 

Water,  solvent  power  of,  40. 

Water,  specific  heat  of,  40. 

Water,  synthesis  of,  37,  41. 

Water,  the  freezing  of,  39. 

Water-vapor  in  the  atmosphere,  108. 

Weights,  relative,  of  the  atoms,  121. 

Welsbach  light,  203. 

White  lead,  309. 

White  vitriol,  260. 

Witherite,  248. 

Wood's  fusible  metal,  182. 

Wrought-iron,  276. 


Xenon,  111. 

Yellow  prussiate  of  potash,  279. 
Yttrium,  271. 

Zeolites,  270. 

Zinc,  258. 

Zinc  amalgam,  258. 

Ziucate  of  potassium,  259. 

Zinc  blende,  258. 

Zinc  carbonate,  260. 

Zinc  chloride,  259. 

Zinc  dust,  258. 

Zinc,  granulated,  258. 

Zinc  hydroxide  soluble  in  sodium  hy- 
droxide, experiment,  264. 

Zinc  oxide  and  hydroxide,  258. 

Zinc  sulphate,  260. 

Zinc  sulphide,  259. 

Zinc  sulphide,  precipitation  of,  ex- 
periment, 264. 

Zinc,  white,  258. 

Zirconium,  207. 


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